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WORKS TRANSLATED BY 
WILLIAM T. HALL 


PUBLISHED BY 
JOHN WILEY & SONS, Inc. 


F. P. TREADWELL’S ANALYTICAL CHEMISTRY 
In Two VOLUMES 
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ANALYTICAL CHEMISTRY 


Based on the German text 
OF 


FB. P. ee Vat, Pu. D., 


Professor of Analytica emistry at the Polytechnic Institute of Ziirich 


TRANSLATED AND REVISED 


. BY 


WILLIAM T. HALL, 8.B., 


Assistant Professor of Analytical Chemistry, Massachusetts Institute of Technology 


Votume I 


QUALITATIVE ANALYSIS 


FOURTH ENGLISH AFTER THE EIGHTH GERMAN EDITION 
TOTAL ISSUE, FIFTEEN THOUSAND 


NEW YORK 
JOHN WILEY & SONS, Ino. 
Lonpon: CHAPMAN & HALL, Limtrep 
> +7 19th. aes ates 





Copyright, 1903, 1906, 1913, 1916 
‘4 Arey et . WIMDIAM 8. HALE os es a 











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PREFACE TO THE FOURTH ENGLISH EDITION 





THe German text upon which this book is based was written by 
an American who has taught for many years at Zurich. The first 
English edition appeared as an authorized translation by one who had 
been teaching analytical chemistry for only three years at that time. 
It was translated largely as a result of a chance remark of his assist- 
ant, R. W. Balcom, who deplored the fact that the students could not 
read German readily enough to make use of the German text as a 
reference book. At that time, the translator was using, as he has 
always used, the excellent book of A. A- Noyes as a laboratory 
manual in Qualitative Analysis and the extremely useful text of 


_H. P. Talbot for the preliminary work in Quantitative Analysis. The 


results obtained by asking the students to purchase both volumes 
of this book in addition have been exceedingly gratifying. Better 
examination papers have resulted and there have been fewer unneces- 


‘sary questions asked in the laboratory. 


Recently Professor Noyes has greatly changed his text on Quali- 
tative Analysis and this has unquestionably had considerable influence 
upon the preparation of the fourth English edition of this. book. It 
has been so thoroughly revised and so largely rewritten that it is no 
longer fair to Professor Treadwell to publish the book as a literal 
translation, although the writer remains in thorough sympathy with 
Professor Treadwell’s views and does not wish, in any way, to dis- 
claim the great benefit and inspiration he has derived from close study 
of the original text. He must, however, express his -obligation to 
other texts, particularly to those of Noyes, Stieglitz, Béttger, and 
Ostwald, from which many of the ideas introduced into this text have 
been copied. The general plan of the book has been kept the same, 
but greater stress has been laid upon the theoretical side of the subject, 
particularly with regard to the applications of the mass action prin- 
ciple, the ionization theory, and the theory of oxidation and reduction. 

The translator wishes to acknowledge his indebtedness to Mr. 
Donald Belcher, who has read all the proofs of this edition and 
offered many valuable suggestions. 

Wiuu1amM T. Hatt. 


MassacHusetts InstiTuTE or TECHNOLOGY, 
February, 1916. 


iii 


frre f) 





PREFACE TO THE FIRST EDITION. 





Havine been repeatedly requested by former pupils to publish 
the lectures on Analytical Chemistry given by me at this Institute 
since 1882, and not having time then to do it myself, I permitted 
the ‘“‘ Verein der Polytechniker ” in 1885 to print in manuscript form 
the notes of one of my students. 

This output met with such a friendly reception that in 1888 a 
second edition became necessary. Subsequently I decided to revise 
the material thoroughly and publish it in book form; this text-book 
of Analytical Chemistry represents, therefore, a somewhat amplified 
repetition of my lectures. 

The book is intended not only for laboratory use, but also for 

self-study. With each element the mineralogical occurrence, crys- 
- talline form, and isomorphous relations are briefly mentioned. Then, 
after explaining the reactions, the methods of separation are given 
in the form of tables; because, contrary to the views of many, I have 
in this way obtained the best results in teaching. These tables are 
summarized charts by which the student can quickly find .his 
bearings. : 

Much weight is placed upon the determination of the sensitive- 
ness of the single reactions, as explained on page 75, because the 
beginner becomes in this way at once familiar with the solubility 
of the most important salts, and also with simple stoichiometrical 
calculations. The approximate solubility of potassium chloroplatin- 
ate, for example, is found from the following determination of the sen- 
sitiveness of the reaction by which it is formed: 

If 100 cc. of the solution contain 0.156 gm. potassium, one finds 

that the formation of the chloroplatinate, at ordinary temperatures, 

only takes place on addition of a little alcohol; but on increasing 

slightly the amount of potassium in the solution, it takes place imme- 

diately. We can, therefore, assume that the solution, which con- 
Vv 


vi PREFACE TO THE FIRST EDITION 


tains 0.156 em. of potassium per 100 cc. water, is saturated with chloro- 
platinate; hence the amount of the latter may be calculated: 


Ke : KePtCle =0.156 : x; 
78.3 : 485.8 =0.156 : a; 
z=0.97. 


The result shows that 100 cc. of water, at ordinary temperatures, 
dissolve 0.97 gm. of K2PtCle, while accurate determinations at 20° C. 
have given the value 1.12. The difference, about 12 per cent, is. 
explained by the facts that we did not work at exactly 20° C., nor 
with absolutely pure water; the solution also contains an excess of 
chloroplatinic acid, whereby the solubility of the potassium chloro- 
platinate is diminished; evidently the values obtained in this way 
permit a very good comparison of the solubilities of the different 
salts. From the sensitiveness of the reaction between a potassium 
salt and tartaric acid, the solubility of the potassium acid tartrate 
may be found to be 0.38; so that the solubility of the potassium 
chloroplatinate is to that of the potassium acid tartrate as 0.97 : 0.38; 
the potassium tartrate is about three times as insoluble as the chloro-- 
platinate, etc. } 

The size of the book does not permit going into the microchemi- 
cal detection of the different elements. We have, however, in the 
excellent work of H. Behrens, ‘ Anleitung zur mikrochemischen 
Analyse,’’ a reference book of the highest rank. 

In publishing this, the first volume of the work, I beg of my col- 
leagues and fellow chemists to kindly inform me of any errors or — 
omissions. 

F. P. TREADWELL. 

Zuricu, April 29, 1899, 





INTERNATIONAL ATomMIc WEIGHTS, 1916 





| 





Symbol | eight Symbol | Atomic 
Aluminium. ...... Al 27.1 Molybdenum.....| Mo 96.0 
Antimony........ Sb 120.2 Neodymium...... Nd 144.3 
PII 6 OS A 30.88. ll Neots cece Ne 20.2 
VAPORS, 82. ac. S52 As 74.96 || Nickel........... Ni 58.68 
DAMM Ss esse vs Ba 137.37 || Nitrogen......... N 14.01 
Pema. yt... Bi 208 .0 Osnaium. 665. 22s Os 190.9 — 
vt Satake Apel B 11.0 Ceyeon Obs O 16.00 
Bromine.......... Br 79.92 || Palladium........ Pd 106.7 
Cadmium. ....... Cd 112.40 || Phosphorus...... P 31.04 
Caesium.:........ Cs 132.81 Platinum........ Pt 195.2 
Calotum: .. ... 5.22; Ca 40.07 || Potassium....... K 39.10 
CeROORy oes, ee C 12.00 || Praseodymium. . Pr 140.9 
SOO a Se ee Ce 140.25 |} Radium: 2.0. 5... Ra 226.0 
Chlorine... ...2..: Cl 35.46 || Rhodium........ Rh 102.9 
Chromium........ Cr 52.0 Rubidium........ Rb 85.45 
IOAN ais. toon oa Co 58.97 Ruthenium....... Ru 101.7 
Columbium....... Cb 93.5 Samarium........ Sa 150.4 
ee A ae lines Cu 63.57 || Scandium........ Se 44.1 
Dysprosium...... Dy 162.5 Selenium......... Se 79.2 
LS 0) 0 a Er 167.7 SINGOR ess. Pa sees Si 28.3 
Europium........ Eu 152.0 SUG aaah ate Sie Ag 107.88 
Fluorine.......... F 19.0 Sodnanhis o> 2s.40 8 Na 23.00 
Gadolinium....... Gd 157.3 Strontium........ Sr 87.63 
SL a ar Ga 69.9 Gaetars séo4% eens S 32.06 
Germanium....... Ge 72.5 Tantalum........ Ta 181.5 
Glucinum?!....... Gl 9.1 Tellurium........ Te 127.5 
OE Rea Ne aA Pa Au 197.2 Terbiim. 34.5... Tb. 159.2 
PIBMUIOS, i505 0. He 4.00 || Thallium........ Tl 204.0 
Hydrogen........ H 1.008 || Thorium......... Th 232.4 
pO a a In 114.8 Then 0 os es Tm 168.5 
BOUMNOs: 05. 6 aie! 6 I 126.92 ‘Tihs Oe foes heen: Sn 118.7 
a a ee Ir 193.1 ‘Pitaniine: S75 a8 Ti 48.1 
LO aa ee eae Fe 55.84 || Tungsten........ W 184.0 
MRerpion <=. <5. Kr 82.92 || Uranium......... U 238 .2 
Lanthanum....... La 139.0 Vanadium....... V 51.0 
OP Re eee arte Pb 207 : 20: Si Menos.: 2. eceeees Xe 130.2 
PAGNUED f 22-3 olera'. Li 6.94 || Ytterbium....... 
Lutecium......... Lu 175.0 (Neoytterbium) . Yb 173.5 
Magnesium..... Mg 24.32 || Yttrium......... Yt 88.7 
Manganese....... Mn $4.93. Zine. sia eee Zn 65.37 
Mercury......... Hg 200.6 Zirconium....... Zr 90.6 























1 Also called Beryllium, Be. 


TABLE OF CONTENTS 


PART I 
GENERAL PRINCIPLES 
PAGE 
I NM Re SE Gol awa ic sie Sy Wadg ts Fas 4 lee tps v Frontispiece 
International Atomic Weights, ROPERS gees es aE fos a 64 aS al RO, XV > 
Qualitative and Quantitative Analysis............... 2c ces ce sce c eee ceeeees 1 
RN ATU UE WU 20 sg ca da dS mro W an 3 ob '6 vod we, wtelbice'b ono Mca RD BTS 1 
emGnEYy OF SuleCtrOly Ue DissOCintion. 2... ise cee bce e ene seas veneees 4 
Mearention. OF Common MICCITOLYteS, . 2 ocak ids co veces csv cceaccscevdnkees 10 
NINE AE CIES ROQEI dea g's cere iho eo die alba 6 a8 ove ere so wR a aa 
Equilibrium between a Solid and a Liquid.................. 2. eee ee ee eee 11 
Chemical Equilibrium and the Mass Action Law.................0.e0ee eee 13 
Equilibrium between a Solid and Two Liquids..............0... 000-20 eee 17 
Influence of Changes in Concentration upon Ionization...................--. 18 
IEE IE REN sg ne Pe Ne Poros AE UES ehie yO ae aie cs Eee eee 19 
NNR CME a eons WA Sg Ge a atohe seek me tak oS i dD cow og aie ob vb Og REE 24 
IMEI NEC TENSES MONEE oe ot cx yar ss ee eA Ae, Ripe ee SEs pb Pane POP Rees 26 
NURI HEIRS EME) PURMSIOTAONY 355)... stein Sh a eG sale Se Po bnle obo eos Fb ke eS 27 
Electromotive Series and Oxidation Potentials.....................0020000- 37 
ET Sab Th AMERTNONY LORE 55 65 css vcs ok ace Ss Wk alee Cae edible O40 40 eee eae 45 
ERR Ne ot iy CN Scr enc oes SaaS OS cd 5 aie nb Ww ale oma ae 48 
MSE EIENISE VOR 65 cea gia satiieg oie Ae wale sales so oben oe sei eee ead 52 
Detection of Acids and Bases. ‘Theory of Indicators....................6.. 54 
IER UNIR EE OMI rary c Micah tw Sing x sags Me 6 boa ee View ed bye Mae ohn ew 56 
Putration and Waehing Precipitates... 0... 6c cee tent cnet ene e ewes 58 
EOOUOIGAL DOMILIONS: nF ee cee ees BRON a URE Ge He Mas Ry Wa K cote Mie oone 58 
ES CTE ROME WE PO, So De yk ge cle Vig rk dao Wp a mg a tiw Ste Mace 0 wie, bo edema 61 
Mareen 15, Ci RO CAIS INUC GETOUIIG, 6 ka cs ek hele Ws bce o Slee sae es aeiseiobe 70 
PN ML TSTTE, 078 SLE nt 55a bain kf hide a Gov mc Sewisc ks 8 ds wd lec bg alereece 71 
Determination of Sensitiveness of Reactions... .........++ Os Wa Sx ca cee ee 
PART II 
REACTIONS OF THE METALS (CATIONS) 
Group V (ALKALIES) 
MINS ere tee gr Ooi al ertN paigd SEs fo a ws, ore e's AV Gee Wowie oe RTO 78 
Re ra Nyy a Vis ee oie gh pic aretaiere Wisk de Mele ad go he LA CER MES 82 
Ce Gl See Ae ALE OG of Fale cas #6 KET Kis WV as Sele pew 87 
SRN NIN As ae oe eb fsb gia be 8 FINO side 6 Oia ie oles Leh PRES View eel n eee Ew Ew 93 
Separation of the Metals of Group Vv eau hE Cab caine Oe db UCR ON ae se wep 96 


x CONTENTS 


Group IV (ALKauine Earrtus) 


PAGE 
General Characteristic Reactions. «.... 6 655 6.3 ek FAG pa eae Bebe vena coeee 101 
ANNU 5 Sa 5 digs creo b ote abide trp ite piece ahate Ga bikes Be neite RIE oak a eae 102 
Peramneiiem | ss ss Sak. Ss Sis de bo a a Hagin 3 me Wee be eee nae a eee 106 
Ne a ne re on ai orem re SN er Cer SiN a ey 108 
Separation of the Metals of Group IV............. cece cececcensvceveveees 110 
Spectoscopic Analysis... .ccvslevvcrcvenseces ei iad» bad 4 Ora ele holon Se 115 

Grovp III 

PE MCUEASITIONIUL 5705 os so a'5. bc Sore 6 a wks aia Bd Ga URE a eh onie ea pee eke ala ae 125 
MSPRPOMYMINATIN Sioa a6 6 cso ssc cen, Kp b OCRig bs oreee W pte Dlahaence Bibb he eeckeadece ok ene 132 
OND ays icc 6 ¥a wb Ue we 6 Wa 0 bck wn e'm Fb URE Sta SPW A cscatener el ehal on Tegal ae 142 
Gi k(t, ae Olen ERAS ERY eh eee SLR Pacem RES Oe 154 
SR AUMTATURIYY 5.5 (a: 5 o-oo Gavin slb0ae’e Bin wih oes am emia sie M ace Sect a secalaile ib lm. 8 oye ee 157 
RMI ATNN sod, 5 cos aia. 0 0:e preikie ace vi wd alnatg ie plata lalate a Be eae ets tn 160 
MIO. 65g 3k ois! obo pcwla.8's Coad we MAS Lat Ae pildd eeelininlecele accuse 172 
MOE MING os 55-5. Oos ia 6; w' 5;0'p, Bin 'v 0 din Whabe mae wid 'e tle Soe Silence beib eciase, dia vs d1atasia a ee 178 
DEES six: ib 'o sia baal oad Wows Wise te OOP ek CE ROS Be eh ea 184 
Separation of Group III from Groups IV and V.............e cece ce eeees 188 


Analysis of Group III in Absence of Phosphates, ...screeeererreverseereess 189 


Grovp IT 
BROROINT on. scion. naw sates si0o.9/08 2 GEL Tb Unie eaiy.s 6 a eho k Wee ka eae ne 195 
MIR a Diao! at roe SR ar a hs cae Oca Ge oes (ba ets ace ae ae 205 
MUORTAININ eo h.6se 8 ed 0 Waele ace 3 a ARS AMO OE dD iano ORE pa ee 210 
oe, PEE ET Te CE LPP T STEEP ey ee 214 
EEE CLEC E PET PPE Es ie PrN Res PL 221 
EP ELO EOE TOL FORE TE EOL ET COPE e PS PET TEL CER Ss ee 223 
MOY. 8S os ks Roe EUS Ra bk 08 CAE ae we a os Wh ka es Daa 242 
Me Ba tio Vein sg sen'eio. 0 V8 ae bes AM Ee hb wid le aM RLR ede Galak Cee nd bie 0 ne 249 
RAMUS IE Gia le s2't ae eb ab a Wu. b-acacedwi'e winie WAS wonlareerel Mmalon alta em aoe bila: Graeme ann 261 
BMORETANTRNY 5-594. a 555s, Dea av s'0 048 ae 0 hoo eivice '0'8 6 Nala oe ENG ad COE oy ee 266 
Separation of Gold from Platinum... 6.5.6.6 scabsesvectevceen.ve scenes emi 270 


Aanivaie OF Group Ics) 9.4 00:9 bs 0b'ee doe 64-0 VEEN 008 SR 271 


Group I 
RUBS Fi ia ksaiw eno se ek niples o chao e's io G Reade o! slilaid ain atau 0 bikin © ana eaeeen a e 278 
Mercury (Mercurous Compounds)... oi. swicsccccceccuvceneneevn eup Oaeee 201 
BROS age keh ae eo civiee'p woin sao Olu ORs nd oN aN aan see lg ee 205 
Aunbvaw of Group 1, + 5. «.> sav adh din he lale's:p’> Vast 4d wlesele 9 0h be ee eae ear 282 
PART III 


REACTIONS OF THE ACID CONSTITUENTS (ANION 8) 


Division of the Acids into Groups... ........sccecececceccuceececsecesencs 284 


CONTENTS xl 


Group I ! 
PAGE 
Hydrochloric Acid............. CW ibe SN ied gm aais ba LRA se bees patwardcw made 285 
EN Si iine a's o\e Ka ad Had hoa ee lwo 0 PO Goctehan A ME Sig? Se PET OY Fo DL TN 2 291 
EDINA COMB ki ai sr ala,g Bie gi wie ble dv.oe once WldAav n'y hoe via 0 wicks SERS tA SneRe 294 
IRIE AUD rent sry een eo aN orig as Hoe gst are S hle a ay hoes emewee Ue 297 
ERs vere CA aL aie Lis tek OE Ue hae bin in aed ee hig Deall epprn eeee 299 
DOOD fare Ae Oy Mars pase AT Rew Ve DOSES kore vibe be aise oe pa 301 
NG Say cee nots PERT ENN UNS AE a i chore nS ociealees * ner he aod 304 
Detection of HCl, HBr, and HI in the presence of one another.............. 306 
Detection of Halogen in the Presence of Cyanide.................0000eeeee 309 
NMED RMR aol gb ara oa) Gene Oras aly alesis ccola oie ' Ad oes eMbined y tiee oh Oeio ER 310 
nh ee SOMA ha aC SVK ow Cds Palo a als AE bias diy Haale eae e slbe 316 
Detection of Nitrogen in Organic Substances.............0.000000 00 cece eens 316 
Detection of HCN in Presence of HCl, HBr, HI,H,[Fe(CN),], Hs[Fe(CN).] and 
MER a SE AN gee Coorg sa fe ig oe iva oi ka wip Aw 0 Sm nike Ai a o/b. wee 4 oianeeaie Bad 317 
EN PM ea aso f, 5 Sec Ovi g-5-47e'e a Se el pete Wine Side a>, WAN ge vie aeRO 318 _ 
IR RAR IE elias tite et Sr Nt | Sd ge alana oak hai we ate etgien 321 
RPM tira eee ER ee ee ik ota ah g SWAG Abe a pasha yO wa ea 323 
Dection of HCNS in Presence of Ha'ogen and Cyanide....................4. 325 
Detection of HCNS, H,[Fe(CN).] and H3[Fe(CN)g]..............0.0 2 eee eee 327 
MTC VANIG ACI. fo. cds vce cwe es ey ew ren Sie WAN ark ea ge at Nis ¢ uh 6 ea 328 
Grovp II 
ENRON CG US eden ok St gius flay Yolk w avy. 9 viet eitus-a Poe Roo ai boa ia. a Rip aps ea 329 
Hydrosulfuric Acid (Hydrogen Sulfide).............. ccc cece cece eee c eens 334 
Re a. coos Ayo SSE RVG ive AUD FE SSRN vie eNO ks CSW pie rainray mse 339 
SERIES IRE ela nel 2 DERE IN ers gh ty SO a Pe om ke ae A 341 
IIR One NM LL a Dag PAI See ere gly samme abe ae elie 343 
IRI FRG tees 5, os vee bie oe wy Pre Gees WO eae € 40S PAE 345 
Grovp III 
I ON sad ad onl slelenns Ga Peete SP OMAR bbe ed RGR ES VU ewoed owes 347 
RTE RES RINGS RIDE SEAL Ae tas i gi Pay. 0 ag Ie De em aE ae 352 
NS 2 Sag na gd arsig Gk GM 9. ST ae RW Lees bee eae rdees 356 
I SS De esa ioral ee hin sak Pelee Fe Sees Foss dC eae fe be bide ben came 357 
TRIE ee Ne te eh RoR ke 5 cole bs cab ace bes wach b sf Wray alate 360 
TERRE ed sare isi a) ae gk as Be ar er waar Ee LL 362 
Meer ait Ie Skates FREE S64 cas > pre Re 8 pint ha bus a epee TTS 366 
UE ON ee Bek pera g 504-4 bain @ Wea le ovialk Mera e melee gale 369 
SD SOCIETIES a ee A ee ee a SO aC eT I ae 372 
TE ONE ES IS eh ek Sa ee Pa a Rare era aT 374 
TO: ASME ov as 'sag pee AP ae ee ater aca wloy'a Da UK Wins ce tea Re ekeTorarals 375 
Group IV : 
EATERIES ANNES es, A cus fav alo Chips 0d hatdws weeds eee Re ee Chae Re EL Oe 377 
MINIS coh, sun a'a Re aih cts SP aan Owed Cob US RENT aeN Cage Cem eM EK RES 382 
PE PRMMIONSN TDN on caks DaP A nce ble Se vwloidly avo 4 6 SNA Sie, s es bie pe ORM e Wb U.N 6 oR ale 224 


IU EE pee tos ola ain, ots bap bie 3 kee CO Fe ME ONCE RY COREE ee Be 229 


xii CONTENTS 


PAGE 
Chrome Acids oo. oe een Cos oss se ROO Ee ea eke bie a Dee 135 
Thinsulfurie Aeid.. . 3.553 on. ah Pawo ke te ee 387 
Detection of Sulfurous and Thiosulfuric Acids in the Presence of Hydrogen 
NEO 5s eo sien slg satan vle’s Chios wiv vie Wino cre aban Stas ae ean gree 390 
Groupe V 
DUAGRID ACIDS Sooo 5ga5 ony ols 08 Sn aeitig Ok ylaicio a ois tw blvd Seas eee 392 
Detection of Nitric Acid in the Presence of Nitrous Acid.................-. 395 
RBIOPIC ACID ook os hp eeaacde g Oa lo pve ls betes we dieses en ee 397 
Detection of Hydrochloric, Nitric, and Chloric Acids in the Presence of One 
PAMNGTNO Fo og oS 5 oy6. vn 0.8 pm WL plc Bn abe wk aoa teas ks 398 
Preeti Se 5 eae es boo oo wm obo wos eae pine ne es 400 
PCPRTIC CIO ACID 00's Sones coe bila bla Su sensi vib aS ehlaep ok bites ay baste Chet 401 
Monopersulfuric Acid (Caro’s Acid). 2 .3..0).7.005 oa. eek oa een spe 403 
Group VI 
SUREE ONO ACHE i Boos said ce bk sg ARS Ep ee pees 404 
BR VAIN UNONES PRI 5 oe ah dks nee PA oon wield aa Gly Pee ee eee a G0 aaa 406 
FLVOTORUOSNCIG. ACI «5.5 554° 85 v4. tyes b ndeise hh wh oc owiehe OME Oe eS ee 411 
Group VII 
PRPS OUE save a'a's bpce'e PSR See 4, 6-bs05G OSG alba BURR Sle a aca Nie Saget ta an 413 
‘Treatment of Insotuble-Bailicsted.. ...... co i a Wee sso baie 9s Oo a 417 
MMR Ss ag on ln oo Sia 0 Ba OP) Blaise eee eetere ean wah a ek dnote apo eae aaa 420 
PART IV 
SYSTEMATIC ANALYSIS 
Analysis of Solid and Non-metallic Substances... .........00 0c eeeeeeeeeues 423 
Prelinimary. Examimation « . 660006556 5.6 a hae 0 5s ve Ba we OE ee 423 
Bomucn of the Substamce. ¢ oo... os ee eo a ee eie es ee eauee ee ae ee 430 
SIAC TIO 5.600 Sie ne ok vind bo cle ho VR A ee Oy bale ge 432 
Methods of Getting Insoluble Substances into Solution. ................+005 436 
Examination for the Metals:(Cations). .....05..60. oso. eo sk nee tie 440 
Examination for the Negative Elements casein, we a Saee wie sells eeepc 446 
Analysis of Alloys...... Sw wb mele SRUe RR Ro Le G5 Salm k cg eaUe Lae cra ee 452 
Ried inh ee FOR aees aera 452 
PART V 


REACTIONS OF SOME OF THE RARER METALS 


Group V (ALKALIES) 


CMAN ea Peis os be oie ah wc bre phe bole Gal TE IONE VO eee oa viel ae 456 
Bea AI iio is ds ec cs oo os So 00s Dee CER fore k «ae 457 
Bt NE Ra nae CN Bean A oe Se dees) 7 eR Sg ee Pe 458 


Detection of Lithium, Rubidium, and Cesium in the Presence of Sodium and 
WNC oo ec en denis oy ARIA aes A REAR e ots aT ae See 459 





le Ven ease a ee 
- ‘ a aw a 


CONTENTS xili 
Group III 

PAGE 

NT teas oa Wis Cs v6 Sha) 8 0S vias’ Perec araiots ieiyze!<-k.s.a/ee Cra aie wate ora oanS 461 
ME etc oP ape eg etl A PRR PWC age keto 2 462 
IN Sear ec hts MEN re Sg ce Tia des x bi co 5 vn age a 6 edn olga cae ce 465 
EE II ae Sa dials. 3-6 ale 0 baa hid Vivian 4.46 Av'o. co oboe Wan bree ors 467 
ED errs ROE Loe iE thay ees tain Uo gw a 30 bs ole view unlobic buen 469 
MN rach ce ae tp Eie Oc isrg @ aul ole/e Ugalns Sis & Lady vd vee de cat ehabwn voles 472 
ME Nn CG vig ho 2 eam ha Oey ods be bh LSG SL ede hab eee 473 
II UES TAMUNEINAILO © 500° 5"5 Gib a < p.o 0:6 0b arco ofc calc Slee cdvicecscudncsvarensibac. 478 
Ih ei tray Pe ee tals Soe ep ain sb a ee ieee Gee Oa Cow eae 480 
Ne Nowe nedee naeeeucaen wae 
Separation of Tantalum from Niobium..................seccceccscceseces 484 

Group II 

ICE DED Riest ye, et ON Ne ey ee ey a de hin eo h,0 0.4 oe aed vas CaaS 485 
Re ae I oe Ee SY ad ae yiare. da ca wigte Neral he Waa Oba wa 487 
MIE re en ens vc PSA Sh to 4x PGA PoNin wie es 6% ha Reine 491 
RT ett fa a gre s-\\clan tna a wlarseh § acece Xb 5, bale ce bre tym es ween ee VES veer * 494 
EE Rin ree et MEET Pe ie 5 ded win il kw acm w sle's GE Oop a 496 
I RR SERS a TROL PPAR Aa ae” Ry Se pilpeo ry Ca RE aE ORE ROD? 499 
CM toa he ee 28 Rb ca oaxe. aie. Pare chee FECL , BACT eS es, 504 
REI De cr Te etre Le ol BLS rm ie me arg ae one Sy SM Ee wig ne Oe Bieid Chee 507 
IE CO Bien SSL, Meant means Sure Gre try sl ajeia We Wie 5.4 o/s’ chawig biog areca Ons 509 
NEN Le as et Cue as Beg Ane Cae e 6 VERE Oe bee 8 Uwe de eke 510 
MER EAT vier STS ood IER ha ly lke MAA Y inky A eb bow a! 0 oo ee eS 512 


penaration of the Platinum Metals... .cvrccsicccsneceeeseveevereqvensecns 516 





QUALITATIVE ANALYSIS 





PART I. GENERAL PRINCIPLES 


By Chemical Analysis is understood all those operations which 
are performed in order to determine the constituents of a chemical 
compound (or a mixture of chemical compounds). Chemical Analysis 
is subdivided into Qualitative Analysis and Quantitative Analysis. 

Qualitative Analysis treats. of the methods for determining the 
nature of the constituents of a substance, while Quantitative Analysis 
treats of the methods for determining in what proportion the con- 
stituents are present in any compound or mixture of compounds. 

In order to recognize a substance we change it, usually with the 
help of another substance of known nature, into a new compound 
which possesses distinctive properties. This transformation we 
call a chemical reaction; and the substance by means of which the 
reaction is brought about, the reagent. 

We distinguish between reactions in the wet way and reactions 
in the dry way. 


I. REACTIONS IN THE WET WAY 


For the purpose of qualitative analysis only such reactions are 
applicable as are easily perceptible to our senses. A reaccion is 
known to take place—(a) by the formation of a precipitate; (6) by 
a change of color; (c) by the evolution of a gas. In other words, 
the sense of sight is used chiefly in qualitative analysis and most of 
the reactions employed are visual ones. The sense of smell also 
aids in identifying many substances. Thus the vapors of hydrogen 
sulfide, hydrogen cyanide, bromine, carbon disulfide and a great 
many other substances have very characteristic odors. Some of these - 
vapors are poisonous, so that in trying the odor it is best to gently 
waft a little of the vapor, by a motion of the hand over the substance 
to be tested, in such a way that the vapor reaches the nostrils greatly 


2 GENERAL PRINCIPLES 


diluted with air. The sense of taste is sometimes useful, but is rarely 
employed on account of the danger of poisonous effects. The sense 
of touch sometimes furnishes a little aid; thus graphite has a peculiar, 
greasy feeling, and paralysis of the tongue or eyelid is temporarily 
imparted by the alkaloid cocaine and certain allied substances. 

When an aqueous solution of barium chloride is mixed with dilute 
sulfuric acid, a white crystalline precipitate of barium sulfate forms: 


BaCle+ H2SO4 = 2HCl1+4 BaSO.. 


A precipitate of identically the same chemical composition can be 
formed from any other soluble barium salt or by using a solution of 
any soluble sulfate instead of sulfuric acid. 

The addition of a little silver nitrate to an aqueous solution of 
barium chloride causes the formation of a white, curdy precipitate 
of silver chloride which darkens on exposure to light: 


BaClo+2AgNO3=Ba(NOs)o+2AgCl. 


The same precipitate is formed when hydrochloric acid or any other 
chloride is used instead of the barium. chloride and when any other 
soluble silver salt is used instead of silver nitrate. 

In the same way there are certain properties which are shown by 
aqueous solutions of all acids. Blue litmus is turned red, carbonates are 
decomposed with effervescence, and metals are dissolved. These so- 
called acid properties are due to the hydrogen of acids, which behaves 
in an essentially different manner than the hydrogen of other com- 
pounds. 

Bases also show certain characteristic reactions which can be 
traced to the hydroxyl, OH, that they contain. An aqueous solution 
of a base turns red litmus blue and reacts with the hydrogen of an acid 
to form water. : 

The aqueous solutions of acids, bases and salts, therefore, show 
reactions which are characteristic not so much of the dissolved sub-_ 
stance as a whole as of its constituents. This is a very important 
point. It enables us to test for the constituents of a solution more 
or less independently of what. other constituents may be present. We 
can test for barium in just the same way whether it is present as 
chloride or as nitrate, and we can test for chlorine by the same reagent 
no matter whether the chlorine was originally present as hydrochloric 
acid or as some other chloride. This is quite remarkable, because the 
chemical properties of a compound are usually quite different from the 
sum of the properties of its constituents. The properties of the chemi- 


‘ 


REACTIONS IN THE WET WAY 3 


cal compound water show little similarity to the properties of either 
hydrogen or oxygen gas. The properties of sodium iodide are quite 
different from those of metallic sodium and of free iodine, and those of 
potassium chlorate are quite distinct from the properties of the potas- 
sium, chlorine and oxygen which it contains. Aqueous solutions of 
acids, bases and salts, however, actually do show additive properties, 
i.e., sodium chloride in solution shows properties which the sodium 
of any other sodium salt will show, plus other properties which any 
other chloride will show. This suggests the hypothesis that the 
aqueous solution of an acid must contain the acid hydrogen, to some 
extent at least, in the same condition as in the aqueous solution of 
any other acid; that an aqueous solution of a base must contain a 
part at least of its hydroxyl in the same condition as the aqueous 
solutions of any other base; and that the metals and non-metals of salts 
must be present in very much the same condition irrespective of the 
nature of the original salt. This would mean that when the acid, 
the base, or the salt is dissolved in water, it is decomposed to some 
extent into smaller units. 

Not only the chemical behavior of aqueous solutions of acids, 
bases and salts indicates that the constituents are present in a con- 
dition such that they may react independently, but also the physical 
behavior of the solutions. The boiling-point of a solution of sugar 
in water is higher and its freezing-point lower than that of pure water. 
It has been found that the rise in boiling-point and lowering of the 
freezing-point is proportional to the number of molecules of dissolved 
substance present. This rule holds so exactly for solutions of organic 
substances dissolved in organic solvents that it serves for the deter- 
mination of molecular weights. When, however, it is attempted to 
determine the molecular weight of an acid, a base, or a salt, by deter- 
mining the boiling-point or the freezing-point of its aqueous solution, 
it is found that the molecular weight thus found is always too small. 
In other words, a study of the boiling-point or freezing-point of acids 
bases and salts indicates that the original molecules of the acid, base 
or salt have been more or less split up into units smaller than the orig- 
inal molecule. : 

Finally the electrical behavior corroborates this view. It is well 
known that substances behave differently toward the electric cur- 
rent; some are conductors of it and others are non-conductors. 
Metals are good conductors and sulfur is a non-conductor. Again, 
the conductors are divided into two classes. Metals belong to the 
first class and conduct electricity without experiencing any change 
except that they become warmer. Conductors of the second class 


4 GENERAL PRINCIPLES 


are chiefly aqueous solutions of acids, bases and salts. Simultaneous 
with the conduction of the current they undergo a chemical change, | 
and decomposition products are obtained at each electrode. 


Theory of Electrolytic Dissociation 


If we insert between the poles of an electric battery a piece of 
rock salt or some pure distilled water, there will be no electric 
current in the circuit; a piece of fine platinum wire placed in the 
circuit will not be made to glow. The solid rock salt and the dis- 
tilled water are non-conductors of electricity; they are non-electro- 
lytes: If, however, we dissolve rock salt in distilled water, and then 
insert the solution between the poles of the electric battery, the 
platinum wire will be brought to a bright glow, showing that the 
salt solution is a good conductor of electricity— it is an electrolyte. It 
is thereby proved that by dissolving the non-conducting rock salt 
in non-conducting water an essential change of the former has taken 
place. We can make the same observation with all acids, bases 
and salts. In an anhydrous state they are non-electrolytes, while 
in aqueous solution,* on the other hand, they are electrolytes. These 
phenomena are readily explained by the theory of electrolytic dis- 
sociation proposed by Arrhenius{ in 1887. According to this 
theory, all electrolytes are partially decomposed in aqueous solution — 
into electrically charged atoms or atom-groups called tons; and the 
extent of this dissociation increases with dilution, until with very 
great dilution it is practically complete. For every degree of dilution 
there exists a certain state of equilibrium between the ions and undis- 
sociated molecules. . 

When the non-electrolyte rock salt is dissolved in water, it breaks 
up, according to the equation 


NaCl @ Na*+Cl- 


into positively charged sodium ions and negatively charged chloride 
ions. { 

All salts, acids, and bases behave like rock salt. Thus sodium 
sulfate decomposes according to the equation 


NaeSO4u @ Nat+Nat+S0O¢-, 


* They are also electrolytes in the fused state. 

t Z. phys. Chem., 1, 631. 

t Many chemists prefer to designate the positive ions by small dots and 
the anions by small inclined dashes. The above equation is then written: 
NaCl = Na+Cl’. 





THEORY OF ELECTROLYTIC DISSOCIATION 5 


fy and sodium hydroxide. into 
NaOH @= Nat+OH-. 


By this theory of electrolytic dissociation the phenomena of elec- 
trolysis may be explained. very simply: If we insert the two poles of a 
scurce of electricity into an electrolyte, one of the poles, the anode, 
is charged with positive electricity, and the other, the cathode, with 
negative electricity. The electro-positive anode repels the electro- 
positive ions (cations) and attracts the electronegative ions (anions); _ 
and the latter, as soon as they come in contact with the anode, give 
up their negative electricity, become neutral and separate out.* The 
same thing happens at the cathode, where the electro-positive ions 
(cations) are discharged. The amounts of electricity which are neu- 
tralized at the electrodes are always renewed by the source of the 
current, so that the process is continuous. 

The electric charge on one atomic weight in grams of a univalent 
ion is 96,500 coulombs ;t on an atomic weight in grams of a biva- 
lent ion the charge is twice as much, and on a trivalent ion three 
times as much. To deposit one atomic weight in grams of silver at 
the cathode, therefore, it is necessary for 96,500 coulombs of elec- 
tricity to pass through the solution, and there will be a simultaneous 
discharge of an equivalent weight of anion at the anode. One 
.coulomb is the quantity of electricity which is represented by the 
flow of 1 ampere for one second, 96,500 coulombs, therefore, represent 
96,500 ampere seconds or 26.8 ampere hours. 

The transport of electricity in aqueous solutions takes place only 
by means of the ions; the undissociated molecules take no part in the 
process. The concentration of the ions and the conductivity of the 
solution are quantities which are proportional to one another. It 
is possible, therefore, to determine the extent to which a solution is 
dissociated into its ions by measuring the electrical conductivity of 
the solution. 

The laws governing electrolysis were well understood by Michael 
Faraday in 1834, and he gave the name of ions to those parts of the 
solution which migrate toward the electrodes (cf. p.-10). The 
positive electrode is called the anode and the negatively charged ion 
which is attracted toward it is called the anion; the negative electrode 
is called the cathode and the positively charged ion which is attracted 





* In many cases the discharged substance at once reacts with the water, forming 
ions again with evolution of either hydrogen or oxygen gas; these gases are, there- 
fore, secondary products of electrolytic action. 

{ This quantity of electricity is called one Faraday. 


6 . GENERAL PRINCIPLES 


toward it is called the cation. In Faraday’s time it was thought 
that the first action of the electric current was to decompose the 
molecules of the substance into the ions. About 1885 Arrhenius 
made the simple observation that all those solutions in which the 
dissolved substances have abnormally low molecular weights, as deter- 
mined by boiling-point elevation, by freezing-point lower'’ng, or by 
some similar method, are solutions which permit the passage of the 
electric current—they are electrolytes; while solutions which give nor- 
_ mal results in the determination of the molecular weight of the dis- 
solved substances are non-electrolytes. Since 1885, therefore, it has been 
believed by most chemists that electrolytic dissociation, or ionization, 
takes place when an acid or a base or a salt dissolves in water, 

This accounts for the fact that aqueous solutions of all silver salts 
show similar reactions. They all contain the silver cation, and the 
silver cation is different from ordinary metallic silver chiefly on account 
of the fact that it bears a large electric charge. Most of the reac- 
tions of qualitative analysis are carried out in aqueous solutions with 
electrolytes. Most of the separations employed and most of the tests 
are by means of reactions which are characteristic of the ions. For 
this reason, a proper understanding of the theory of electrolytic dis- 
sociation is necessary in the study of qualitative analysis. 

Let us interpret the action of a dilute solution of an acid upon a 
dilute solution of a base. In a dilute solution of hydrochloric acid, 
for example, the hydrogen chloride is almost completely ionized, and in 
a dilute solution of sodium hydroxide the base is also almost com- 
pletely ionized. Hydrochloric acid and sodium hydroxide react 
together to form water, which is itself but very slightly ionized. The 
reaction between the dilute solutions of ila ys acid and sodium 
hydroxide may be written: 


H*+Cl~+Nat+OH~ =H20+Nat+Cl-. 


By subtracting the ions which appear on each side of the equality 


sign, the equation becomes: 
H*t++OH~ =H.0O. 


According to this, the neutralization of a dilute solution of an acid by 
a base is merely the reaction of hydrogen ions with hydroxyl] ions to 
form undissociated water. This is known to be true, because if the 
reaction takes place with 1 gm. of hydrogen and 17 gms. of hydroxy] the 
heat evolved is 13,700 calories. This same amount of heat is evolved 
when an equivalent amount of a dilute solution of hydrochloric acid 
is neutralized by a dilute solution of potassium hydroxide, or when 
the hydrochloric acid is replaced by another acid such as nitric acid; 


THEORY OF ELECTROLYTIC DISSOCIATION t 


it represents merely the heat of formation of a molecular weight in 
grams (one mole) of water from hydrogen ions and hydroxy] ions. 

Similarly it can be shown that when an acid acts on a metal with 
the formation of a salt and liberation of hydrogen gas, the quantity 
of heat which is developed depends only on the nature of the metal 
and is independent of the acid. The anion of the acid really does 
not take part in the reaction at all. 

The main assumptions of the Arrhenius theory of electrolytic dis- 
sociation are as follows:. When an acid, a base or a salt dissolves in 
water its molecules are immediately dissociated to some extent into 
smaller fragments called ions. These ions are charged with electricity 
and the sum of the positive charges residing on the cations is exactly 
equal to the sum of the negative charges residing upon the anions 
and the whole solution is electrically neutral. The dissociation is a 
reversible reaction and all electrolytes may be considered to be com- 
pletely ionized at infinite dilution. Except for the dependence 
resulting from the electrical charges and the consequent attractions 
and repulsions between ions, the ions may be regarded as independent 
constituents with individual and specific chemical and physical proper- 
‘ties. When a substance dissolves in water and is only partly dis- 
sociated, then when the ions are removed, either by electrolysis or 
as a result of*chemical reaction, the substance will at once dissociate 
again to form new ions. 

While it is true that nearly all acids, bases and salts are ionogens, 
yet the extent to which the ionization takes place when the substance 
is dissolved in water varies greatly. Thus a molecular weight in 
grams of hydrochloric acid dissolved in 10 liters of water yields about 
seventy times as many hydrogen ions as an equivalent quantity of 
acetic acid; a similar comparison can be made with regard to sodium 
hydroxide solution and ammonium hydroxide. In round numbers, 
hydrochloric acid is about seventy times as strong an acid as acetic 
acid and sodium hydroxide or potassium hydroxide is nearly seventy 
times as strong a base as ammonium hydroxide. 

On the other hand, a molecular weight of acetic acid will neutralize 
the same weight of sodium hydroxide that a molecular weight of hydro- 
chloric acid does, and a molecular weight of ammonium hydroxide will 
neutralize the same weight of acid that a molecular weight of sodium 
hydroxide does. In a solution of sodium hydroxide and of hydro- 
chloric acid of the above concentration the original molecules are 
about 90 per cent ionized, and when the acid and alkali are mixed 
the principal change is the union of hydrogen ions and hydroxyl ions 
to form water. When acetic acid of the same concentration is used, 


8 GENERAL PRINCIPLES 


there is present at the start only 1.3 per cent of all the hydrogen in the 
form of ions. These ions will at once react with hydroxyl ions 
to form water, but there is always a tendency for the acetic acid to 
dissociate, and when the ions disappear as fast as they are formed 
the ionization continues and soon all of the molecules of acetic acid 
will have dissociated. In the neutralization of acetic acid with 
sodium hydroxide, the final heat effect will not be simply that of the 
union of hydrogen ions with hydroxyl ions, but will also involve the 
energy required to cause the acetic acid to dissociate. When a sub- 
stance ionizes as soon as it dissolves, the heat effect of ionization cannot 
be distinguished easily from the heat of solution. Just as some sub- 
stances dissolve with absorption of heat and some with evolution of 
heat, so it is found that the ionization process may likewise be asso- 
ciated with either an absorption or evolution of heat. 

It is interesting to note, and this is a matter of considerable 
importance, that the salts of weak acids and of weak bases are usually 
ionized nearly as much as the salts of strong acids or of strong bases. 

When a bivalent acid dissolves in water, the two hydrogen atoms 
do not dissociate to an equal extent. The ionization takes place in 
two stages. Thus with sulfuric acid the first stage takes place in the 
sense of the equation: ' | 


H2804 @ H*t+HS0O,-. 


The fact that the reaction does not necessarily take place completely 
is indicated by using the double arrow sign instead of the equality 
sign. When the above reaction stops there is a state of equilibrium 
between the three substances H2SO4, H* and HSO4~. The HSO,- 
undergoes a secondary dissociation as follows: 


HSOs. = Ht+S0.-. 


The extent to which these reactions takes place depends entirely 
‘upon the dilution. If half a molecular weight in grams of sulfuric 
acid is dissolved in 10 liters of water, the primary dissociation will 
take place to about 90 per cent of the entire quantity of acid present 
and the secondary dissociation to less than 50 per cent. If the solution 
is extremely dilute, both reactions will take place almost completely. 

In the case of carbonic acid, the primary stage 


H2CO3 @ Ht+HCO;- 


ordinarily takes place only to a fraction of 1 per cent and the second 
stage 
HCO3 @Ht+CO3= 


THEORY OF ELECTROLYTIC DISSOCIATION 9 


to an inappreciable extent (cf. p. 10). With hydrogen sulfide the 
relations are similar. 

On the other hand, the salts of these weak acids will dissociate 
almost completely as follows: 

NazCO3 = Nat+Nat+C03* 
NazS = Nat+Nat-+S-*. 

The fact that the extent to which a substance ionizes is dependent 
upon the concentration of the solution has already been indicated 
(p. 4) and will be demonstrated mathematically a little. farther on 
(p. 18). The concentration of a solution shows the quantity of dis- 
solved substance present in a unit of volume, and the numerical 
value representing this concentration depends entirely upon the 
units in which the mass of the dissolved substance and the volume 
of the solution are expressed. If w represents the mass of dis- 
solved substance in grams and v is the volume of the solution 


expressed in liters, then = is the concentration of the solution 


in grams per liter. Concentrations are often expressed in these 
units, but the weight is not the most convenient unit for ex- 
pressing the mass, especially in matters of theoretical discussion. 
A solution containing 1 gm. of dissolved substance A is rarely equiv- 
alent to one containing the same weight of a substance B. It is much 
more convenient to measure the mass of the dissolved substance in 
terms of the number of molecules present. According to the sug- 
gestion of Ostwald, the term mole.has been given to the molecular 
weight of a substance in grams, and when the concentration of a solu- 
tion is expressed in the number of moles present in a liter, the so- 
called molal concentration is obtained. The objection still remains, 
however, that one molecule of a substance A (e.g., hydrochloric acid) 
is not always equivalent to one molecule of a substance B (e.g., sul- 
furic acid). To overcome this difficulty, concentrations are often 
expressed in gram equivalents per liter, using the univalent substances 
as the standard. Thus a mole of hydrochloric acid is one equiva- 
lent weight in grams and half a mole of sulfuric acid is one equiv- 
alent weight in grams. A solution which contains one equivalent 
weight in grams of dissolved substance is called a normal solution; 
one containing two equivalents in a liter is a twice-normal solution; 
and one containing half an equivalent is a half-normal solution. All 
things considered, this is the best way of expressing concentrations. 
The following table will be found useful in studying the disso- 
ciation of electrolytes. It gives the approximate percentage ioniza- 
tion of substances present in 0.1 N solution at 25°. In the case 


10 GENERAL PRINCIPLES 


of polybasic acids, the value opposite the formula of the acid shows 
the fraction of the whole molecule which undergoes the primary 
dissociation into one hydrogen ion, that opposite an ion with a 
univalent charge shows the extent to which this ion undergoes a 
secondary dissociation, and that opposite an ion with a bivalent charge 
shows the extent to which it undergoes a tertiary decomposition, — 
forming a third hydrogen ion from the original neutral molecule of the 
acid, 


- Ionization Values of Common Electrolytes* 


In 0.1 normal solution 


Per Cent 
Salts of the type BTA (ez, KINO) Sea ae eee 84 
Salts of the type B,+A= or B* t+A,7 (e.g., KoSO4 or BaCl)f............. 73 
Salts of the type B;stA=, or Bt + +A;- (e.g., K3Fe(CN)s or AICI;)........ 65 
Saltsiof the type BttA™ (e.@.. MgSO)... os ates sce so bs ap eee 65 
Moyne Ss hy Ooh Ee SSeS ee 90 
Ba(OH). ees EB syhtoly “eine Sie 6 fb. pb apauw Bele te Sues te wpe ah etihts Shae 5] Cue: e, lene, sees mea rae ee ae 80 
18 | Sa a ee ae eee ote es ace MM ee Mica Ma vr a 1.3 
HCl, HBr, HI, HSCN, HNO;, HCIO;, HCO, HSO,, HsCrO............ 90 
H3PO,, H;AsQOu, H.2SOs, H.2C20,, HSO,— a) dhe betel gah oer oe ae BRL aE rae et Ea 20-45 
MNT OSS i ee Ce Ct ae 7-9 
HC.H;02, HC.0,, PSO ge ee nr eh ie Se ee eee 1-2 
HS, H.COs, HPO, = ’ HCrO,~ 4L 5. Mellen fe fa hails Noe ASS RRA Snel SRS Ae hes) De te ica eens ee 0.1-0.2 
PCa IA e, THON, BOO o.oo b sene as) <a Nica ae eee 0.002-0.008 
Rs MES Fes te ee ee ise Reo cS a 00001-0002 
SPR OTE BG BD alice Gh bee ihc Ha 8 6 obebay SARE wie AOR Fae RN oe ee se ere 0 .Q000002 


Nomenclature of the Ions 


As sieaeas mentioned (p. 5) Faraday in 1834 was the first to use 
the words ion, cathode, cation, electrode, anode, etc. ‘These names are 
all derived from Greek roots. Faraday’s idea was that the electricity 
entered the solution at the positive pole and passed down to the 
negative pole. The word zon is the Greek word for wanderer or traveler 
spelt in Latin letters. Anode is from the Greek dvé (ana) up; 600s 
(odos) a path. Cathode is from the Greek xara (kata) down; 600s 
(odos) a path. The two electrodes are considered as the doors or 
paths by which the current enters and passes out of the solution. 

Two methods are in common use for designating the ions. Thus 
the ions of hydrochloric acid are designated as H* and Cl” or as H’ 
and Cl’. Small plus and minus signs are used in this book rather than 





*From A. A. Noyes, Qualitative Chemical Analysis. 
t Exceptions: CdCl, ionizes to about 47 per cent, HgCl to about 0. 01 per 
cent, and HgBr,, HgI., and Hg(CN)s less than HgCh. 


y = 


eve eh os 


NOMENCLATURE OF THE IONS 11 


the dots and dashes simply because this is the present practice in 
the journals published by the American Chemical Society. Many 
writers prefer to use the other system because it takes up less room; 
in the case of the polyvalent ions the use of the plus and minus 
signs is often very cumbersome. 

Purely as a matter of convenience, an attempt has been made 
to devise a system of rational nomenclature for theions. This method 
has been adopted in a number of excellent text books but it is not 


in common use. According to this system, the names of the cations 


are obtained by adding the termination -ion to the stem of the name 
of the corresponding metal, using the Latin name whenever possible. 
When a substance forms several ions differing from one another only 
in valence, the names of such ions are designated by Greek prefixes 
indicating the number of charges residing on the ion. The names of 
the anions are derived from the names of the salts. If the name of 
the salt ends in -ate, these last three letters are replaced by the ending 
-anion, except in the case of the carbonate ion, which is called carbanion. 
The names of anions from salts ending in -ite are formed by replacing 
these three letters with -osion. The anions from salts whose names 
end in -ide are obtained by replacing these letters with the ending 
-~idion. 'The hydrogen ion is called hydrion and the hydroxy! ion is 
called hydroxidion. The following table illustrates the use of this 
system which was proposed by Walker: 3 


NAMES OF CERTAIN IONS 























Symbol. Name. aces reg Symbol. Name. Anion of 
Agt Argention Silver Cl- Chloridion Chlorides 
Cat + Calcion Calcium ClO— | Hypochlor- Hypochlorites 
Cut + Dicuprion Cupriec copper osion 
Fet + Diferrion Ferrous iron ClO,;— | Chlorosion Chlorites 
Fet++ | Triferrion Ferric iron ClO;—_ | Chloranion Chlorates 
Ht Hydrion Hydrogen ClOg | Perchloranion | Perchlorates 

(Acids) 33 Sulfidion Sulfides 
Kt Kalion Potassium SO;= Sulfosion Sulfites 
Nat Natrion Sodium SO Sulfanion Sulfates 
NHgt Ammonion | Ammonium NO;— _ | Nitranion Nitrates 
OH Hydroxidion | Hydroxides 
(Bases) 








Equilibrium between a Solid and a Liquid 


Most of the reactions used in analytical chemistry involve either 
the solution or the precipitation of some substance. It is important, 


12 GENERAL PRINCIPLES 


therefore, to consider briefly the. relations which exist between a solid 
and its solution. : 

Potassium nitrate on being brought into contact with water at once 
begins to dissolve. The rate of solution is influenced somewhat by 
the amount of surface exposed by the salt, a fine powder dissolving 
more rapidly than a single large crystal. At first the substance dis- 
solves quite rapidly, particularly if the liquid is kept stirred, but 
gradually the speed slackens and finally a time comes when the water 
at. a given temperature will dissolve no more of the salt. The solu- 
tion is then said to be saturated with the salt and it makes no differ- 
ence how much potassium nitrate is available in excess of the amount 
required to form a saturated solution, the solution when once satu- 
rated at any temperature will dissolve no more salt. 

The quantity of salt required to form a saturated solution varies 
with the temperature, more so with potassium nitrate than with 
many other salts. At 0° the saturated solution contains only 1.3 moles 
of potassium nitrate, whereas 2.7 moles dissolve at 20° and ” moles 
dissolve at 100°. 

If a solution of 5 moles potassium nitrate is prepared by dis- 
solving the salt in hot water and the solution is then cooled to 20°, 
we obtain what is called a supersaturated solution. A state of super- 
saturation can be maintained for some time provided care is taken 
not to disturb the solution in any way. If the supersaturated solu- 
tion is agitated, or, better, if a tiny fragment of potassium nitrate is 
thrown into it, crystallization starts and continues until finally the 
solution only contains 2.7 moles of the salt, which is the quantity of 
potassium nitrate required to form a saturated solution at 20°. 

The solubility of a substance at any temperature is usually deter- 
mined by two methods: first, by shaking up the salt with water until 
a saturated solution is obtained; second, by forming a supersaturated 
solution and allowing the excess of the salt to crystallize out. Usually 
the values obtained by the former method are a little lower than the 
values obtained by the latter method; a slightly undersaturated solu- 
tion is obtained in one case and a slightly supersaturated one in the 
other. 

When a solution of potassium nitrate is brought into contact with 
more of the salt, whether more of the salt will dissolve or not is 
determined solely by the concentration of the solution. If it is 
saturated already with potassium nitrate, no more of the salt will 
dissolve; if unsaturated, more salt will dissolve to form a saturated 
solution. The equilibrium between a liquid and a solid which is 
soluble in it is determined solely by the concentration of the solution. 


CHEMICAL EQUILIBRIUM AND THE MASS ACTION LAW 13 


The absolute quantity of substance and the absolute quantity of solu- 
tion have no effect upon the final equilibrium. 

Chemists prefer to look upon a state of equilibrium as a condition 
of dynamic equilibrium rather than as one of static equilibrium. 
Instead of thinking of the saturated solution of potassium nitrate as 
one which has no tendency to dissolve more potassium nitrate, it is 
preferable to consider the solution as one in which the tendency to 
precipitate potassium nitrate is exactly balanced by the tendency 
to dissolve potassium nitrate. When the solution is undersaturated 
and more salt is available, the tendency to dissolve is greater than the 
tendency to precipitate and when the solution is supersaturated the 
tendency to precipitate is greater than the tendency to dissolve. 

The equilibrium principle is the same in the case of difficultly 
‘soluble substances. It requires only 0.0015 gram (=0.01 millimole) 
of silver chloride to form a saturated solution in water. If more 
than this quantity of silver chloride is produced as a result of a 
chemical reaction taking place in an aqueous solution, all the excess » 
silver chloride will be precipitated. The solubility is so slight that 
the precipitation is practically complete. 


Chemical Equilibrium and the Mass Action Law 


If hydrogen sulfide gas is passed into a solution containing zinc 
chloride, a white precipitate of zinc sulfide is formed: 


ZnClz+He2S = ZnS+2HCl. 


If the precipitate of zinc sulfide is filtered off and treated with hydro- 
chloric acid it will dissolve: 


ZnS +2HCl1=ZnCl2+H28. 


Similarly, the addition of ammonium carbonate to a solution of 
calcium chloride in water causes the formation of a white precipitate 
of calcium carbonate: 


CaCl2+ (NH4)2C03 =CaCO3+2NH4Cl. 


The precipitate can be dissolved, however, by boiling it with ammonium 
chloride solution. 

In each of the above cases, there are evidently two opposing 
tendencies—the tendency of zine sulfide to precipitate and the tend- 
ency of zinc sulfide to dissolve; the tendency of calcium carbonate 
to precipitate and the tendency of calcium carbonate to dissolve. To 
express the fact that the reaction may go in either direction it is cus- 


14 GENERAL PRINCIPLES 


tomary to writé the symbols separated by a double arrow instead of 
by an equality sign (cf. p. 8): 


ZnCle+H28 = ZnS+2HCl, 
CaCle+(NH4)2CO3 @ CaCO3 +2NH.,Cl. 


Such reactions are called reversible. It was once thought that 
reversible reactions were of rare occurrence, but it is now customary 
to consider all chemical reactions as reversible, although in many 
cases and especially in most reactions used in analytical chemistry, 
the reaction goes so completely in one direction that only a negligible — 
quantity of one or more of the initial substances remains unchanged. 
In general, when two substances A and B react with one another at a 
constant temperature to form C and D, then, to some extent at least, 
C and D react to form A and B, and equilibrium is reached when the 
ratio of the product of the concentrations of A and B to the product 
of the concentrations of C and D has a definite, constant value. This 
value is characteristic of the equilibrium between the compounds 
involved. | ; 

In the above case, the reaction may be expressed as follows: 


A+B 2C+4D, 


in which A, B, C and D represent four different substances reacting 
in the molecular proportions indicated by their symbols. The condi- 
tions of final equilibrium is expressed by the mathematical equation: 


[A] X[B] 
[C]X[D] 


in which [A], [|B], [C], and [D] represent the final concentrations of the 
four reacting substances and k is some definite number called the 
equilibrium constant. The value k varies with the temperature. 

If more than one molecule of substance takes part in the reaction, 
the conditions are somewhat more complicated. This is expressed by 
the general equation 


=k, 


mA+nB = pC+¢qD 
and the final equilibrium is expressed mathematically 
[A] x[B]” 
[C}? x[D]* 
in which [A], [B], [C], and [D] represent, as before, the concentrations 
when equilibrium is reached. 





=k, 


ee a ee a, eee eee See 


CHEMICAL EQUILIBRIUM AND THE MASS ACTION LAW 15 


This is the so-called law of mass action, which was discovered by 


-Guldberg and Waage in 1867. It is to be noted that it is the concen- 


trations, or masses present in a unit of volume, rather than the actual 
masses of the substances, which find expression in this law. 
This law apphes to a state of homogeneous equilibrium. A homo- 


geneous system is one in which every part of it is like every other 


part. A mixture of two solid substances is not homogeneous. A 


solution, on the other hand, is homogeneous when it is thoroughly 


mixed, as it is impossible to distinguish any difference between differ- 
ent portions of the solution. Similarly a mixture of gases represents 
a homogeneous system. Such homogeneous systems are called 
phases. A mixture of a solid, a so‘ution and a gas represents three 
phases; two solids, two phases; two immiscible liquids, two phases. 

In the case of the reactions between zine chloride and hydrogen 
sulfide and between calcium chloride and ammonium carbonate a 
precipitate was formed in each case. The mass-action law applied 
only to the zine sulfide and to the calcium carbonate that remained in 
solution. The fact that these substances are only very slightly soluble 
in water favors the progress of the reaction in the direction by which 
these substances are formed. The mass-action law shovs that when 
the concentration of any substance participating in a chemical reac- 
tion is increased, this tends to increase the tendency for the reaction 
to take place in the direction by which this substance is decomposed; 
when any substance formed by means of a chemical reaction is removed, 
this increases the tendency for the reaction to proceed in the direction 
by which this substance is formed. The formation of a precipitate 
or the escape of a gas, as fast as the substance is formed by means of 
a chemical reaction, tends to make the reaction take place more com- 
pletely. If the gas is all boiled off the reaction by which it is formed 
will take place completely. Similarly if any precipitate were absolutely 
insoluble in water, the reactions by which this substance is formed 
would take place completely. 

This law of mass action embodies one of the most important 
principles utilized in analytical chemistry. It enables one to under- 
stand why most of the reactions take place and to establish conditions 
under which these reactions will occur to the best advantage. The 
law has been verified by a great many quantitative as well as 
qualitative experiments. It has been studied, for example, in 
connection with the formation and decomposition of phosphorus 
pentachloride. 

When chlorine gas reacts with cold phosphorus trichloride, the solid 
pentachloride is formed; but if this substance is heated, it breaks 


16 GENERAL PRINCIPLES 


down into its constituents. The reaction is reversible and may be 
expressed as follows: | 
PCls @ PCls3+Cle. 


At any given temperature an equilibrium exists which can be expressed 
mathematically, according to the mass-action law, 
[PCl3] X[Cle] a 
[PCls] : 





in which [PCl3], [Cle] and [PCls5] represent the concentrations at the 
time when equilibrium has been reached. - 

If we desire to volatilize phosphorus pentachloride so that the 
least possible dissociation will take place, the above equation shows 
us how this may he brought about. 

If either [PCls] or [Cle] be increased, then in order that the value 
of the fraction 

[PCls] [Cle] 
[PCls] 





remains constant, it is evident that the concentration [PCl5] must 
become greater; or, in other words, the dissociation of the penta- 
chloride becomes less and there will be practically no dissociation if 
the pentachloride is volatilized in an atmosphere of phosphorus tri- 
chloride or of chlorine. In this way, Wurz obtained for the density of 
phosphorus pentachloride 6.80-—7.42, instead of the calculated value 7.2. 

At Stassfurt the mineral carnallite (MgCle- KC]-6H2O) occurs and 
was evidently formed by precipitation from solutions containing the 
chlorides of magnesium and potassium. This double salt is less soluble 
than pure magnesium chloride and more soluble than pure potassium 
chloride. If carnallite is dissolved in water and the solution allowed 
to evaporate until crystals are deposited, it will be found that the 
crystals consist of potassium chloride. When the carnallite dissolves 
in water, the double salt is decomposed, more or less completely accord- 
ing to the dilution, 

MgCle: KCl @ MgCle+ KCl, 


and for every concentration the equation holds: 


[MgClo] X[ KCl] 


iMeClo aon: COnStane 





If we wish to recrystallize the carnallite, the breaking down of the double 
salt must be prevented as much as possible, and to do this it is merely 
necessary to add an excess of MgCle. As a matter of fact, the mineral 


EQUILIBRIUM BETWEEN A SOLID AND TWO LIQUIDS 17 


is recrystallized at Stassfurt from a 23 per cent solution of magnesium 


chloride. 

This law of mass-action applies to all cases of chemical equilibrium 
that takes place in a homogeneous phase, i.e., it can be applied to 
all reactions which take place between gases and to all reactions that 
take place in solution. 

The law of mass action applies to the equilibrium between an ion- 
ogen and its ions. In this connection, the law is of particular impor- 
tance in the study of analytical chemistry. 


Equilibrium between a Solid and Two Liquids 


Although water, either pure or containing dissolved acid, is the 
solvent most used in analytical chemistry, it often happens that a 
substance is more soluble in some other liquid. Thus free iodine 
is about 400 times as soluble in carbon disulfide as it is in water. - 
When iodine is in contact with both carbon disulfide and water, and 
these two liquids are only slightly soluble in one another, it will dissolve 
chiefly in the carbon disulfide. Moreover, if an aqueous solution of 
iodine is shaken with carbon disulfide, the latter, when it separates 
out beneath the water, will contain nearly all of the iodine. A state of 
equilibrium then exists between the solution of iodine in water and the 
solution of iodine in carbon disulfide. Such an equilibrium is governed 
by the so-called distribution law or law of partition. If C4 represents 
the concentration of a substance in a solvent A and Cz is its con- 
centration in a solvent B, equilibrium is reached when 


a. 

C7 
This is the mathematical expression of the distribution law. The con- 
stant, k, is called the distribution coefficient. In this simple form, 
it is important to note that the law holds only when each concentra- 
tion is expressed in terms of the same molecular species. Thus if a 
substance is dissociated to a large extent in one solvent and scarcely 
at all another, the concentrations involved must be those of the un- 
dissociated salt in each case. It is quite common to find that the 
ions of a substance are much more soluble in water than in any other 
solvent whereas for the undissociated substance the relations are 
reversed. Iodine dissolves to a greater extent in a solution of potas- 
sium iodide than it does in pure water, owing to the formation of KIg3. 
In such a solution the following equilibrium exists: 


KI+I, @ KIsz. 


k. 


18 . GENERAL PRINCIPLES 


If such a solution is shaken with carbon disulfide, the distribution 
law holds only for the free iodine held in solution as such in each 
solution. 

For iodine in pure water and iodine in carbon disulfide, the dis- 
tribution coefficient is z4y. Theoretically it is impossible to. remove 
all the iodine from water by shaking with carbon disulfide, but if the 
carbon disulfide is removed, with the aid of a separatory funnel, and the 
aqueous solution is shaken with fresh carbon disulfide, it is evident 
that the quantity of iodine remaining with the water is negligible, - 
or can be made so by repeating the operation. 

Sometimes in testing for the halides it is desirable to remove free 
halogen from the aqueous solution; to accomplish this, the distribu- 
tion principle is utilized. Ferric chloride is much more soluble in 
ether and hydrochloric acid than it is in water and hydrochloric acid; 
to detect the minor constituents of iron or steel, a large sample of the 
- original material is taken and the ferric chloride removed by shaking 
the hydrochloric acid solution with ether. Perchromic acid is more 
soluble in ether than it is in water; by shaking the dilute aqueous 
solution with a little ether, a concentrated solution in the latter is 
obtained and the presence of the chromium shown by the beautiful 
blue color. 


Influence of Changes in Concentration upon the Ionization of 
! Electrolytes 


If we assume 1 gm. molecule of a weak electrolyte, such as ammo- 
nium hydroxide, to be dissolved in »v liters of solution, the original 
substance will be partly dissociated according to the equation 


NH,OH @ NH4y*+0H- 


into ammonium and hydroxyl ions. If a gram molecule of the base 
is dissociated in the sense of the above equation, then the undisso- 
ciated part will amount to 1—a. 
The concentrations per liter are 
| undissociated dissociated 
NH40H @ NH4t+OH- 

Se. 

Sea ® v 
and according to the mass-action law 








we ae a eae hese aa ay é- 
- 7 se iy Me 


SOLUBILITY PRODUCT 19 


The constant k is known as the tonzation- or affinity-constant; 


_ it is independent of the dilution and is characteristic for every elec- 


trolyte. The expression shows, however, that by increasing v, or 
diluting, the fraction of the molecule dissociated (a) will be made 
larger. 

If to the solution of the base we add n additional ammonium ions, 
by adding solid ammonium chloride, then if is considerably larger 


than a the degree of dissociation of the base will be greatly diminished, 


namely, from a to a1, a value which can be readily computed as follows: 
In the solution there is present per liter 


undissociated dissociated 
NH4OH @ NH4t+0OH— 
l-—a, ai+n a1 
v v v’ 
therefore 
ps (aitn)a1 
: (1 —ay)v : 
If k and n are known, a; can be computed: 


or (n+vk) + V (n+vk)?+40k 
za ; ; 


In the case of 0.1 N ammonia solution the ammonium hydroxide 
is dissociated only to an extent of 1.32 per cent, the dissociation con- 
stant being 0.000018. If we add to 10 liters of this ammonia solution 
2 gram-molecules of ammonium chloride (107.08 gms.), then since the 
ammonium chloride is 93 per cent dissociated at this dilution, we are 
adding 2X0.93=1.86 NHa4 ions. If this value be inserted in the 
above equation, ai becomes 0.00009; in other words, the dissociation 
of the ammonium hydroxide is diminished by the addition of the am- 
monium chloride from 1.32 to 0.009 per cent. The solution now 
contains so few hydroxyl ions that it will not cause precipitation in 
solutions of magnesium salts (cf. pp. 46, 94). 

Similarly the dissociation of weak acids is lessened by the addi- 
tion of their salts. In the case of the stronger acids and bases, the 
effect of adding a neutral salt to the solution is not so remarkable, 
because the stronger acids and bases are dissociated to about the 
same extent as their salts. 








Solubility Product 


Silver chloride is slightly soluble in water; 0.00001 mole (1.5 
mgs.) of the solid dissolves in 1 liter of the solvent. When water 


20 . GENERAL PRINCIPLES 


is placed in contact with an excess of silver chloride, a state of 
equilibrium is soon reached between the solid and the solution. If 
more than this quantity of dissolved substance is present at any time, 
the solution is supersaturated and tends to precipitate silver chloride; 
if less, then more silver chloride will be dissolved. When equilibrium 
is reached the tendency of the salt to precipitate is equal to the tend- 
ency of the salt to dissolve. 

The dissolved silver chloride also exists in a state of equilibrium 
between the ionogen and the ions. This equilibrium apparently 
takes place almost instantly, whereas the equilibrium between the 
solid and the solution is established more slowly. For all concen- 
trations of such a slightly soluble substance as silver chloride, it is fair 
to assume that the mass action law holds rigidly. Applied to the 
reaction 

AgCl @ Agt+Cl- 


and denoting the concentration of non-ionized silver chloride as [AgCl], 
of silver ions as [Ag], and of chlorine ions as [Cl], the law is expressed 
as follows: 


[Ag] X[Cl] 


[AgCl] =f 


In this equation k has a definite value, called the ionization constant, 
which varies with the temperature but has otherwise a definite value 
for every substance. When the solution is saturated with silver 
chloride the value of both numerator and denominator has reached 
the saturation value. If the value of [AgCl] ismade greater than 
corresponds to this saturation value, the solution is supersaturated 
and is no longer in equilibrium with undissolved solid. The value © 
[AgCl] could be used for expressing the solubility of the substance, 
but it is often more convenient to use the ion concentration product 
[Ag] X[CI] which is called the solubility product. In general, if the sub- 
stance A,,B, ionizes into mA and nB ions, the solubility product, 
Sy, is found by the following equation: 


= [A]”x [B}” .< k[|A,B,l, 


in which k is the ionization constant and the concentrations are those 
of a saturated solution. 

Experience has shown that the conditions are somewhat more 
complicated in concentrated solutions such as are obtained with the 
very soluble substances. In future discussion, therefore, the solu- 
bility will be expressed, as a rule, in terms of the solubility product 





a ee es ae o = ee 


SOLUBILITY PRODUCT | 21 


; only when the substance does not dissolve to a greater extent than 


0.01 mole per liter. The table on page 10 shows that binary salts 








of the type represented by AgCl are ionized to about 84 per cent 
in 0.1 N solution, and we have seen on page 19 that the ioniza- 
tion increases as the solution is diluted. In the case of such a dilute 
solution as that of silver chloride (0.00001 normal) the ionization 
is nearly 100 per cent. It is therefore logical, in such cases, to express 
the solubility in terms of the ions, whereas in the case of the very solu- 
ble substances it is better to measure the solubility in terms of the 
mass of dissolved substance. | . 

The numerical values of the ionization constant and of the solubility 
product depend upon the units used in measuring the mass of the 
dissolved substance or ions and in measuring the volume of the solu- 
tion. It is customary, in this case, to express the concentration in 
moles per liter (cf. p. 9). In the following table the solubility of the 
substance is expressed in three ways: in grams of dissolved substance 
per liter, in moles per liter and, finally, in terms of the solubility prod- 
uct, using moles per liter. The solubility of most of the substances 
given in the table is so slight that the quantity dissolved is negligible 
for most purposes. Whenever the word insoluble is used in this book 
it is with the understood limitation that no substance is absolutely 
insoluble in water. 

SOLUBILITIES AND SOLUBILITY PRODUCTS AT ROOM 
TEMPERATURE 

Substance. eames, a Fc pompscaty SoA Solubility Product. 

JS eee 1.1X10-4 | 5.9X10-7 | [Ag]x[Br]=3.5x10-" 
Ag.(CN)s...... 4.3X10-5 1.6X10-7 | [Ag]x[Ag(CN).|=2.6X10-"4 
AgCNS....... 1.4x10-4 8.4X10-7 [Ag] X[CNS] =7.1x10-# 
an 1.5X10-3 1.1X10-5 | [Ag]x[Cl]=1.2 10-1 
AgeCrO........ 2.5X10-? 7.5X10-5 [Ag]? [CrO,] =1.7x10-™" 
AgeCrO7...... 8.3xX10-? 1910-4 [Ag]? X[Cr2O7] =2.7x10-"™ 
7 ee 2.1X10-2 | 9.0x10-5 | [Ag]x[OH]=1.9X10-8 
Oo SEs 3.010-8 1.3X10-8 | [Ag]x[I]=1.7x10-16 
ST: ae 4.4X10-2 1.5X10-* | [Ag]x[I03]=2.310-8 
MPO we... 6.5xX10-3 | 1.6X10-5 | [Ag]*x[PO.J=1.8X10-18 
MIMO. 5 8.0 2.6X10-2. | [Ag]?<[SO.]=7.0X10-5 
ES Se eae 1.3X10-? 4.3X10-5 [Ba] X[CO3] =1.9 10-9 
BaCrQy....... 3.8X10-* | 1.5xX10-5 | [Ba]x[CrOJ=2.3x10-% 
DOE. Cis es 5. 2.5X10- L.1x<io-6 [Ba] x [SO] =1.2 10-10 
PR ge OY isc <8 1.3 7.5X10-3 [Ba] X[F]?=1-7 10-6 
2S 6S At lena 1.3X10-? Loxi0-* [Ca] X[CO;] =1.7 X10-8 
COsG,04;. 2... 5: 8.0X10-3 6.2x10-5 [Ca] X[C,04] =3.8 X10-9 
TOUTS. ds 23. 1.5X<10-! [Ca] X[CrO,.] =2.3 10-2 
Sa ae a 1.6X10-? 2.0x10-4 [Ca] X[F]?=3.4x10-1! 














22 


GENERAL PRINCIPLES 


SOLUBILITIES AND SOLUBILITY PRODUCTS AT ROOM 
TEMPERATURE—Continued 





Solubility in 











Substance. Grams per Liter. Afcles wer Lite: Solubility Products. 
a 9 SOR 2.0 1.5X10-2 | [Ca]x[SO.]=2.3 10-4 
Cus(CNS)2.....| 5.0*10-4 2.1X10-§ | [Cu] xX[CNS]=1.7x10-" 
Cush: so 1.2x10-3 6.0X10-4 | [Cu] x[Cl]=1.410-8 
i Se 3.0X10-4 8.0X10-7 | (Cu) x[I]=2.6x10-” 
oe ae 8.8X10-% 9.2x10-%3 | [Cu]x[S]=8.5x10-45 
a OE AES 3 8.6X10-%8 6.0X10-15 | [Cd] x[S]=3.6x10-% 
YO) 5 8 Rs MRCS Nee Carb, Mr Sace ici tu) as [Fe] X[OH]?=1.6<10-" 
ae 3.4X10-8 3.9X10-19 | [Fe] x[S]=1.5x10-19 
LTO UE S PAROMRNS RAND Pte WR amen Vee ae eR aot [Fe] x[OH]*=1.110-%6 
Hg.Bre....<..- 3.9X10-5 6.9X10-8 | [Hg.]x[Br]?=1.3 10-7 
PS os See es 3.8X10-4 8.0X10-7 | [Hge] X[Cl]?=2.0 10-18 
7 2 Fe eae 2.0X10-7 3.1X10-1° | [Hgo] x[I]?=1.2 10-78 
BROCE eee ens. ata eds eee peas [Hg] X[OH]?=4.3 X10- 16? 
cS ARR etnies ag CMR en cA aE Rom BR {Hg] x[S]=4.0 10-53 
KaPtCle,.: es 11. 2.3xX10-2 | [K]?x[PtCle] =4.9 X10-5 
MgCos...'.\... 4.3X10-! 5.1X10-% | [Mg] X[COs] =2.6 10-5 
Mg(OH)>.\.: .. 1.2Xx10-2 2.0X10-4 | ([Mg]x[OH]?=3.4x10- 
MgNH,PO,,. 8.6X10-3 6.3xX10-5 | [Mg] X[NH4] X[PO,] =2.510-# 
DOOR oes Heck een eae ae See [Mn] X[OH]?=4.0 10-4 
Muse 2. 38X10 +s 3.8X10-8 | [Mn]x[S]=1.4x10-% 
i ee eat ie OKO 1.2X10-12 | [Ni]xX[S]=1.410-% 
(ei ag a 9.7 2.7X10-2 | [Pb] X[Br]?=7.9 10-5 
Pate aS 11. 3.9X11-2 | [Pb] x[Cl]?=2.4x10-4 
PHOO ee 1.1X10-3 4.1X10-§ | [Pb]x[CO,]=1.7x10-" 
Phin ees 6.8x10-! 1.5X10-* | | (Pb) xX[f?=1.4x10-t a5 
Pb;(PO,)2...... 1.4xX10-4 1.7X10-7 | [Pb]*[PO.J?=1.5 10-3 
eS eA eae ae 4.2X10-2 1.5X10-4 | [Pb]x[SO.] =2.3 10-8 
PbCrO,....... 4.3X10-5 1.3X10-7 | [Pb] x[CrO.]=1.8 10-4 
Mi. OG o28 4.9X10-” 2.0X10-" | [Pb] x[S]=4.2x10-%8 
Lt ae 6.6X10-2 3.8X10-4 | [Sr] x[C,0.]=1.410-? 
BroUs..fi3.. 1.0x10-? 6.8X10-5 | [Sr}x[CO3] =4.6 10-9 
ate Ee tas Ns 2.3X10-1 1.8xX10-% | (Sr]x[F]?=2.5x10-9 
BRE os cet 1.1X10-3 6.0X10-4 | [Sr] x[SO.]=3.610-7 
FEBS aa sarees 4.8X10-! 1.7X10-% | [Tl]<[Br] =2.9x10-§ 
oy Cama 3.2 1.2x10-2 | [TI]X[CNS]=1.4x10-4 
=: AR Ae oad 3.4 1.4x10-2 | [Tl] x[Cl]=2.010-4 
<4 Ra git hor 6.4x10-? 1.9x10-4 | [TI]x{I]=3.6x10-8 
MET CS § SERRA ts Ag: Cadet eae ak ay Leah seal Nees [Zn] X[OH]?=1.8X10-"4 
aaa Ri TE 3.5X10-12 | [Zn] x[S]=1.210-28 


a: 10—% 








The above table is prepared from many sources, and the values 
are based, in some cases, upon solubility determinations by methods 
which are now considered inaccurate. 


The table gives a good idea, 
however, of the relative order of magnitude. 


For copper sulfide, 


the table states that 8.8x10-?! gms. dissolve in 1 liter of water. 
Obviously, the experimental determination of such a small value is 


fraught with difficulty and the probable error is large. 


For most 


” 


Sa sl! oA ae -" a a, AS La es ee ili or ts 5 Sy 


SOLUBILITY PRODUCT 23 


purposes such a value represents a negligible quantity and the state- 
ment is often made that copper sulfide is insoluble in water. It is 
instructive, however, to compare the solubility products of the various 


sulfides and important methods of separation have been based upon 
‘such studies. Only two significant figures have been given in the 


table, although it is obvious that more would be justifiable in the 
case of the more soluble substances, while even the first figure is doubt- 
ful for the very insoluble substances. The values are affected to 
different degrees by changes in temperature and the presence of 
other substances in solution. A careful, critical study of all the experi- 
mental data would be necessary to give the proper number of signifi- 
cant figures and it would be necessary to give the exact temperature. 

In computing the solubility products, the assumption has been 
made that the ionization is complete. Such an assumption is not 
permissible with a substance such as ferric hydroxide, and in such 
cases only the approximate value of the solubility product is given; 
the molar solubility and the grams per liter are not stated. In other 
cases the ionization is abnormal as noted. 

Two examples will be given to illustrate the method of computing 
the molar solubility, S,,, and the solubility product, S,. A saturated 
solution of silver iodide contains 3.010-§ gms. (=0.0030 mg.) 
per liter. The molecular weight of silver iodide is 234.8. The satu- 


3.0X 10-6 | es 
S35 = 1-3X10-8. At this 


dilution the dissolved silver iodide can be assumed to be completely 
ionized: 


rated solution, therefore, contains 


Agl @ Agt+I-, 


and since 1 mole of silver iodide furnishes 1 mole of silver ions and 1 
mole of iodine ions, it is evident that the solubility product, S,, is for 
Paw (li=[1.3 X10-*| X[1.3 10-8] =1.7 X10-%=§,. 

A saturated solution of silver phosphate contains 6.510% gms. 
(=6.5 mgs.) per liter. The substance is much more soluble in water 
than silver iodide, but its solubility product is smaller. The molecular 
weight of silver phosphate is 418.7. The saturated solution, there- 


_-3 
2X7 =1.6x10-5 moles of silver phosphate which 


can be assumed to be completely ionized: 
AgsPO4 = 3Agt+PO., 


1 mole of silver phosphate yielding 3 moles of silver and 1 mole of 
phosphate ions. The solubility product is 


[Ag]? X[PO4] =[3 X1.6 X10-5]* X[1.6 X 10-5] = 1.8 X 10-18 =§,. 


fore, contains 


24 GENERAL PRINCIPLES 


Complex Ions 


Silver chloride is slightly soluble in water; 1 liter dissolves about 
1.5 mgs. It dissolves very readily in dilute ammonia. The following 
reaction takes place: 


AgCl+2NH3 @ Ag(NHs3)2Cl. 


A study of the properties of this new substance shows that it dis- 
sociates in aqueous solution chiefly in this way: 


AgNH;Cl 2 Ag(NH3)2*-+Cr-. 


The ionic changes involved in the last two equations may be 
expressed thus: 
Agt+2NHz3 = [Ag(NHs)2]* 


and, in accordance with the law of mass action, the greater the con- 
centration of the ammonia, the greater the extent to which the reac- 
tion takes place in the direction left to right. In a normal solution 
of ammonia, the ratio of the concentration of the [Ag(NH3)e] ion to 
that of the simple Ag ion is about 107: 1. The [Ag(NHs3)o| ion differs 
from the simple Ag ion in much the same manner as the ClO, ClOsz, or 
ClO3 ion differs from the simple Cl ion. It is called a complex cation. 

When potassium cyanide is added to silver nitrate solution a white 
precipitate of silver cyanide is formed: 


KCN+AgNOz3 = KNO3+AgCN, 
but if an excess of potassium cyanide is used the precipitate dissolves 
AgCN+KCN = KAg(CN)z. 
In this case the ionic changes may be expressed as follows: 
Agt+CN7- @ AgCN, 
AgCN+CN™ @ [Ag(CN)e]~ 


and the silver has become a part of the anion. In this case the value 
of the ratio of complex ion to simple ion is even larger than in the case 
of the silver ammonia cation. 

Similarly, when insoluble ferrous cyanide is treated with an excess 
of potassium cyanide, it dissolves, forming potassium ferrocyanide, 


Fe(CN)2+4KCN = K4[Fe(CN)ol. 


This salt gives none of the ordinary reactions of ferrous ions. The 
iron forms an integral part of the complex ferrocyanide ion which has 


COMPLEX IONS 25 


its own characteristic reactions, and during electrolysis alwavs migrates 
toward the anode: 


K,[Fe(CN)¢] — 4Kt+Fe(CN),~~. 


It is, in fact, quite common to find that simple salts, particularly 
in concentrated solutions, are capable of forming such complex com- 
pounds. The simple ions can unite with neutral molecules, or with 
ions of opposite charge, to form complex ions. If a simple ion adds 
to itself a neutral molecule, such as H2O, H2O02, NHs or organic radicals, 
then neither the original valence nor the electric charge is changed. 
Thus the trivalent cobalt ion is capable of forming a deep red ion with 
6 molecules of ammonia and this complex ion has a trivalent charge 
like that of the original simple ion. 


Cottt+6NH3 @ [Co(NH3)¢]t**. 


If, in such a complex ion, one or more of the ammonia groups is 
replaced by a negative univalent ion, the valence of the complex ion 
is reduced one for each atom of negative ion thus entering into the 
complex. : 

(NOz)2|*, (NO2)s3 

In the presence of potassium cyanide, the trivalent cobalt ion 
unites with six CN ions to form a complex which has, in accordance 
with the above rule, a triple negative charge. The valence or electric 
character of a complex ion is the algebraic sum of the valencies or elec- 
tric charges of the constituents. 

As already indicated, the stability of these complex ions varies. 
When the complex is-very stable, the common reactions of the con- 
stituents are not shown. | 

Besides these complex ions certain double salts are known. Thus 
potassium and aluminium sulfates crystallize together, forming an 
alum, KeSO4-Ale(SO4)3-24H20. When this salt is dissolved in 
water, the solution shows all the reactions for potassium, aluminium 
and sulfate ions, and there is little evidence of the formation of a 
complex ion. To determine, in a given case, whether a substance 
is a double salt or a complex salt, it is customary merely to see whether 
the characteristic reactions of the simple ions are shown. A salt 
exists which has the symbol KCr(C204)2:5H20. An aqueous solu- 
tion of this salt readily shows the reactions for the potassium ions, 
but reacts sluggishly when tested for chromium cations or for oxalate 
anions. Evidently the chromium and the oxalate have united to form 


26 GENERAL PRINCIPLES 


a complex anion with a negative valence of one, but this complex is 
not as stable as some of the others that have been mentioned. It is 
probable that there is no sharp distinction between the double salts 
and the complex salts- and probably the double salts are most logic- 
ally to be classed as complex salts of which the complex ion is not 
very stable. As a general rule, those complex salts which are composed 
of neutral salts of strong acids yield complex anions which are largely 
dissociated into simple ions in dilute solution. On the other hand, 
complex ions composed of positive ions and anions of weak acids are 
usually very stable. 


Reactions of the Ions 


As already indicated, most of the reactions used in qualitative 
analysis involve reactions between ions. We have seen.that, in prin- 
ciple, all reactions are reversible and have learned to understand 
some of the laws which govern these reversible reactions. In analyti- 
cal chemistry, it is necessary for the most part to employ reactions 
which take place almost completely in the desired direction. Unless 
a reaction can be made to go nearly to completion in a given direction, 
it is of little value either as a sensitive test or for furnishing a method 
of separation. The useful reactions of qualitative analysis, namely 
those which-apparently go to completion, may be brought into four 
classes: : 

(1) Reactions in which a gas is formed. 

(2) Reactions in which a precipitate is formed. 

(3) Reactions in which a non-ionized substance is formed. 

(4) Reactions of oxidation and reduction. 

When a gas is formed as a result of a chemical reaction and the 
gas escapes, the reaction will go to completion. All gases can be 
boiled out of solution and thus all reactions of this type can be made 
to go to completion. The reaction can be stopped by preventing 
the escape of the gas; this shows that the reaction is inherently a 
reversible one. 

Whenever a substance which has a very small solubility product 
is formed by means of a chemical reaction, the greater part of the sub- 
stance will leave the solution in the form of a precipitate and the 
reaction will go practically to completion. The table on page 21 
shows that the saturated solution of silver chloride contains only 
about one hundred-thousandth of a mole (=0.01 millimole) of solid 
salt per liter. The table also shows that when the product obtained 
by multiplying the concentration of the silver ions by the product 
of the concentration of the chlorine ions in any aqueous solution is 


OXIDATION AND REDUCTION 27 


equal to 1.3X10-!° the solution is saturated with silver chloride. 
By adding an excess of chlorine ions to a solution containing silver ions, 
it is possible, therefore, to precipitate nearly all of the silver. It is 
evident that it will take less silver ions to give the solubility product 
when an excess of chlorine ions is used than is necessary when pure 
silver chloride is dissolved in water. 

The precipitated silver chloride will dissolve completely in potas- 
sium cyanide, because the silver ion forms with the cyanide ion a 
complex which is ionized to such a slight extent that the solubility 
product of silver chloride is no longer reached, even although all the 
chlorine is present in the ionic condition. 

The formation of a non-ionized substance also causes a reaction to 
go to completion. The table on page 10 shows the ionization values 
of a few common substances. This table may be used exactly like that 
of the solubility products to enable one to predict whether a reaction 
is likely to go in a given direction. The equilibrium between water 
and its ions H and OH has been discussed on page 6. The same 
reasoning may be applied to the equilibrium between any other slightly 
ionized substance and its ions; whenever the ions are added sepa- 
rately to a solution, some of the non-ionized substance is at once 
formed. ‘Thus when any acid is added to the solution of a sulfide 
a reaction takes place, partly because the hydrogen sulfide is a very 
weak electrolyte and partly because the substance is a gas. Sim- 
ilarly calcium phosphate dissolves in hydrochloric acid because 
more PO, ions are formed from dissolved calcium phosphate than 
are formed from HzPO,4~ ions in the presence of an excess of H* ions 
from the hydrochloric acid; the reaction takes place because of the 
formation of a non-ionized substance. 

Finally, many reactions of oxidation and reduction take place nearly 
to absolute completion, although all these reactions can be shown ~ 
to be inherently of a reversible type. To understand such equilibria, 
however, it is necessary to discuss oxidation and reduction at greater 
length. 


Oxidation and Reduction 


The term oxidation, in its narrowest sense, signifies the taking 
up of oxygen by an element or compound. Thus ferrous oxide, on 
being heated in the air, is converted into ferric oxide and the reaction 
is called an oxidation. Since, however, ferric chloride bears the same 
relation to ferrous chloride that ferric oxide bears to ferrous oxide, 
it is customary to call the change of ferrous chloride into ferric chloride 
an oxidation, although it is not necessary to think that oxygen takes 


28 GENERAL PRINCIPLES 


part in the reaction at all. This is an interesting example of a word 
in common use which has come to mean a great deal more than it 
originally meant. Indeed, chemists have departed so far from the 
original meaning of oxidation that sometimes the word seems inappro- 
priate, and the use of another word, such as adduction, has been sug- 
gested. Reduction is the exact opposite to oxidation, and whenever 
one substance is oxidized some other substance is reduced. Hydrogen 
was formerly considered to be the typical reducing agent, so that the 
definition for oxidation used to read something like this: Oxidation 
is the addition of oxygen (or its equivalent) to an element or com- 
pound or the taking away of hydrogen (or its equivalent). 
The reaction between ferrous chloride and chlorine: 


2FeClo+Cle = 2FeCls, 
expressed in terms of the ionic theory becomes, 
| 2Fet+++Cle=2Fet + ++42Cl- 


In other words, the diferrion has been converted to triferrion and the 
neutral chlorine molecule has become changed to negatively charged 
chloride ions. In all other reactions in which a ferrous salt is oxi- 
dized, the valence of the iron is increased one, and the modern concep- 
tion of oxidation and reduction is summed up very simply as follows: 

Oxidation is the increase in the valence of an element or radical in the 
positive direction; reduction is the increase in the valence of an element 
or radical in the negative direction. Oxidation involves the assumption 
of positive charges or the loss of negative charges and reduction in- 
volves the loss of positive charges or the assumption of negative charges. 

According to the electronic conception of the constitution of matter, 
the atom of an element consists of positively charged corpuscles and 
negatively charged corpuscles or electrons. The mass associated 
with the positive electricity is much larger than the mass associated 
with the equal charge of negative electricity. The number of positive 
and negative electrons which make up the atom is probably a very 
small multiple of the atomic weight of the element. ‘The mass asso- 
ciated with a unit negative charge is so small that it may easily be lost, 
but only under conditions such that it is accepted by some other atom. 
The originally neutral atom which loses the electron thus becomes 
positively charged and the atom which accepts the negatively charged 
electron becomes negatively charged, and a tube of force holds the 
two elements together in a so-called chemical compound. In the light 
of the electron theory, therefore, an element is oxidized when it loses 
an electron and an element is reduced when it receives an electron. 


OXIDATION AND REDUCTION 29 


This is the simplest, and at the same time most comprehensive theory 
of oxidation that has ever been suggested. | 

‘Oxidation, according to this conception, is essentially an electric 
phenomenon. This theory suggests the thought that it ought to 
be possible to accomplish oxidation and reduction simply by means of 
electric energy. As a matter of fact probably every oxidation and — 
reduction can be brought about in the electrolytic cell if the proper 
conditions be maintained. Using the conventional symbol @ to 
designate a unit charge of positive electricity and ©) to designate a 
unit charge of negative electricity, but bearing in mind that the nega- 
tive electricity is alone transferred and that the only way an element 
can gain in positive charge is by losing one or more negative electrons, 
we may express oxidations in the electrolytic cell as follows: 


Fett+ @—Fettt or Fet*—©— Fet tt. 


Such oxidations take place at the electrode called the anode. Con- 
versely, at the cathode, ferric salts can be reduced to the ferrous 
condition: 


Fet*+++@©-— Fett. 


_ Not only may all oxidations and reductions be accomplished with 
the aid of the electric current, but, vice versa, an electric current may 
be produced by a proper arrangement of the components of any reac- 
tion of oxidation and reduction. Thus some ferric chloride and sodium 
chloride solution in a small beaker may be connected with a second 
beaker containing sodium chloride by means of a U-tube filled with 
dilute salt solution. If a platinum electrode is placed in each beaker 
and the terminals are connected with a sensitive voltmeter, no cur- 
rent will pass through the wire. On pouring some hydrogen sul- 
fide water into the beaker containing sodium chloride, a decided deflec- 
tion of the voltmeter needle is at once observed, showing the passage 
of an electric current. The negative current enters the voltmeter 
from the solution containing the hydrogen sulfide and passes on to 
the ferric chloride solution and back, through the salt-bridge, to the 
hydrogen sulfide solution. The chemical reaction that takes place is, 


2Fet t+49= — 2Fett+8. 


The electric current is produced as a result of the oxidation of the 
sulfide ions and reduction of the ferric ions. 

Oxidation and reduction reactions are inherently reversible reac- 
tions, like all other chemical reactions, and are effected by the concen- 


30 GENERAL PRINCIPLES 


trations of the reacting substances. Thus, in the above experiment 
the intensity of the electric current can be greatly increased by using 
a soluble sulfide instead of hydrogen sulfide, the former being more 
largely dissociated and yielding a larger concentration of sulfide ions. 
Or, by adding a fluoride to the solution of ferric chloride, a fairly stable 
complex ion, Fel’, is formed and the current slackens, owing to the 
decreased concentration of the ferric ions. 

The most important oxidizing agents used in analytical chem- 
istry are the halogens, nitric acid, potassium permanganate, potassium 
dichromate and hydrogen peroxide. 

The most important reducing agents are nascent hydrogen and 
metals, sulfurous acid, hydrogen sulfide, stannous chloride and hydri- 
odic acid. 

The oxidizing action of halogens depends upon the conversion 
of the neutral halogen into halogen anions. 

The oxidizing action of halogen upon ferrous ions results in the © 
formation of ferric ions and of halide ions: 


2Fet*++Cle > 2Fettt+2Cr, 
or 
2FeCle+Cle = 2FeCls. 


The action of halogen upon hydrogen sulfide is interesting. First 
of all, the sulfid-ion is oxidized to free sulfur, 


H2S+Bre — 2HBr+5, 


but, if the bromine is fairly concentrated, the reaction may go farther 
and the sulfur be converted into sulfuric acid, the whole reaction being 


H28+4Bre+4H20 = H2804+8H Br. 


It will be noticed that it is very easy to balance equations of 
oxidation and reduction by noticing the change in valence. In this 
last equation sulfur goes from a negative valence of two to a positive 
valence of six, making an algebraic change of eight, which corresponds 
to the loss of eight electrons by the sulfur atom. 

The oxidizing action of nitric acid depends upon the reduction 
of the nitrogen. The extent of the reduction depends upon the con- 
centration of the nitric acid and the nature of the substance oxidized. 
The more concentrated the nitric acid, the less it is reduced; the more 
concentrated the reducing agent and the stronger its reducing power 
the greater the reduction of the nitric acid. Nitric oxide, NO, is 
commonly formed, but often other products such as nitrogen peroxide, 
nitrous oxide, nitrogen and even ammonia are produced. 


5 


OXIDATION AND REDUCTION 31 
\ 

In nitric acid, the nitrogen atom has five positive electric charges 
residing upon it. When it is reduced to nitric oxide, NO, it has only 
two positive charges, the nitrogen having accepted three electrons. 
The reaction between a ferrous salt and nitric acid is, 


3Fet*++NO37+4H*3F et +*+NO+2H20, 
or 
6FeSO4 + 2HNO3 +3H2S04 =3Fee (SO) 3+ 2NO +4H20. 


The reaction is doubled in the last instance simply to get an even 
number of molecules of Fee2(SO4)3. The addition of sulfuric acid is 
unnecessary, as the nitric acid can also act as an acid, in which case 
a mixture of ferric sulfate and nitrate is formed: 


3FeSO4+4HNOsz = Fe2(SO4)3-+Fe(NO3z)3 +-NO+2H20. 


The action of nitric acid on a sulfide is interesting. If the nitric 
acid solution is cold and dilute (0.8 N) there is hardly any oxidation 
of the sulfur: 

MnS+2HNO3 = Mn (NOs)2 +H>S. 


If the nitric acid is more concentrated (e.g., 2 N) and the solu- 
tion is heated, the sulfide is changed to nitrate and free sulfur is 
formed. Thus for the reaction between copper sulfide and hot nitric 
acid, each atom of copper requires two nitrate ions, each atom of sul- 
fur loses two electrons and, in accomplishing the oxidation, each atom 
of nitrogen gains three electrons. The reaction may be expressed 
thus: 


3CuS +8HNO3 =3Cu(NO3)2+4H20+35+2NO, 
or 
3CuS+8H*t+2NO3— — 3Cut*+4H20+38S+2NO. 


If the nitric acid is very concentrated, the greater part of it is 
reduced only to NOzg and the sulfur is oxidized to sulfuric acid. The 
reaction may then be written: 


CuS+8HNO3=CuSO04+8N02+4H20. 


The oxidizing action of permanganate depends upon the readi- 
ness with which the manganese is converted into a manganese com- 
pound of lower valence. In the permanganate anion, MnO, the 
manganese has a positive valence of seven. Ordinarily, in acid solu- 
tion, the permanganate is reduced to bivalent manganese cation, 


32 GENERAL PRINCIPLES 


corresponding to a loss of five positive charges, or acceptance of five 
electrons: 
MnO4~+5Fet*++8H*=Mnt*t-+5Fet*+t+4H.20; 


2MnO4 +5H2S+6H* =2Mn*t *+5S+8H20; 


++ 
2Mn0O4~+5S8n* sae 16Ht a 2Mn* F4 58nt *+8H2O; 
2Mn0O4 +10HI+ 6Ht = 2Mntt-+5lo +8Ho20. 


The oxidizing action of a chromate or dichromate ordinarily 
depends upon the formation of trivalent chromic ions. In the chro- 
mate and dichromate ions the chromium atom has a positive valence 
of six, so that for each atom of chromium the reduction corresponds to 
a loss of three positive charges, or gain of three electrons. Potassium 
chromate in acid solution is in equilibrium with the dichromate: 


9CrO4-+2Ht — 2HCrO4~ — H2O+Cr207-. 


In balancing equations, therefore, it makes little difference whether 
we start with the chromate or dichromate, except with respect to 
the quantity of acid required: 


Cr207~+6Fet*+14Ht — 2Crt*+*++6Fet **++7H20; 
Cre07~+38"+14H* > 2Cr++*++38+7H20; | 


+4 
Cro077+38n* ++14H*t > 2Crt*+*+3Snt *+7H20; 
CreO7_ +61 + 14H*t — 2Cr+*+*+31.+-7H20. 


Hydrogen peroxide acts both as an oxidizing agent and as a 
reducing agent. It oxidizes ferrous chloride to ferric chloride and 
it is capable of reducing permanganate to manganous salt. This 
anomalous behavior has been the cause of considerable discussion 
in the literature. It is unnecessary to go into the details of such a 
discussion, but a simple explanation of this behavior will be suggested. 
In all the other compounds of hydrogen and oxygen that we shall 
study, hydrogen has a positive valence of one and oxygen a negative 
valence of two. With hydrogen peroxide, also, it is best to assume 
that the hydrogen has its normal valence corresponding to one posi- 
tive charge. ‘Two structural formulas are at once suggested for hydro- 
gen peroxide, 

H—-O 
OO or Pe 
H—O 


Hy, 
H 


OXIDATION AND REDUCTION 33 


Each of these formulas gives to one atom of oxygen its normal negative 
charge of two units of electricity and each gives to one atom of oxygen 
an equal number of positive and negative charges. It is unnecessary, 
therefore, to attempt to decide which of these formulas is the more 
appropriate. 

The characteristic behavior of hydrogen peroxide may be traced 


to the presence of the atom of oxygen which has an equal number of 


positive and negative charges. In alkaline solution, hydrogen 
peroxide decomposes spontaneously and oxygen is evolved. This 
spontaneous decomposition, with the formation of neutral oxygen, is 
easy to understand on the basis of the assumption that the original 
molecule already contains the atom of oxygen in a very unstable con- 


dition of neutrality. 


The oxidizing power of hydrogen peroxide is due to this atom of 
oxygen. In acid solutions ferrous iron is converted by it to the ferric 
condition. The peculiar atom of oxygen loses its positive charge and 
receives in its place a negative charge; the total change corresponding 
to the acceptance of two electrons: 


2Fet ++-H20.+2H+t > 2Fe++++2H.20. 


The reducing power of hydrogen peroxide also depends upon the 
presence of this atom of oxygen. When in contact with a powerful 
oxidizing agent, such as permanganate or another peroxide; a reac- 
tion takes place and oxygen gas is evolved. It has always been 
assumed, since the classic experiments of Schénbein, that half of the 
evolved oxygen comes from the oxidizer and half from the hydrogen 
peroxide. It is simplest to assume, therefore, that the oxygen is 
momentarily changed to an atom with two positive charges which at 
once unites with negatively charged oxygen in the oxidizer; or, this 
oxygen in the hydrogen peroxide may unite with similarly charged 
oxygen in another peroxide. 


2MnO4 +5H20e2 +6Ht —? 2Mn* *4+8H20+50p; 
Mn0O2+H202+2H* — Mnt *42H20+0Og; 
Coo0s +HeOe +4H* ae 2Cot +43H.O +Ood. 


The characteristic action of the more important oxidizing agents 
has now been considered briefly and it remains to describe the char- 
acteristic behavior of the important reducing agents. Since every 
oxidation involves a simultaneous reduction, all of the above reac- 
tions can be used to illustrate reduction as well as oxidation. 

The reducing action of nascent hydrogen and of metals depends 
upon the conversion of the neutral hydrogen or metal into positively 


34 GENERAL PRINCIPLES 


charged cations. Such a reduction may take place in acid, alkaline, — 
or neutral solution. 
(a) In acid solution, by the-employment of zinc, ete.: - 


Zn+H2S804= ZnSO4+ He 
or 
Zn-+2H*t — Zn* t+Hbp. 


This reaction in itself represents both an oxidation and a reduction, 
inasmuch as the,metallic zinc, which is electrically neutral, becomes 
changed into zinc with two positive charges and, on the other hand, 
the hydrogen in sulfuric acid has lost two charges and become electric- 
ally neutral hydrogen. 


Wy 


Now this nascent hydrogen, as fast as it is formed, may act as a 
reducing agent and serve for effecting the reduction of some other 
element, e.g., ferric chloride; in which case the final changes are the 
oxidation of the zinc and the reduction of the iron: 


Fet t*+4-7n=Zn' T+-2F ett. 


In such cases it is doubtful whether the reduction process really goes 

through the stage of forming nascent. hydrogen. : 
By the action of zine and very dilute sulfuric acid, it is easy to 

transform silver chloride into metallic silver: 


2Agt+Zn — Znt++2Ag, 
or an arsenite into arsine: 


AsO3- +3Zn+9H* — 3Zn**+3H20-+-AsHs. 


The result of this last equation is the oxidation of zine from the metallic 
condition to the bivalent state and the reduction of the arsenic, which 
is given three negative charges in place of the three positive valences 
that each arsenic atom has in AsOg3.. Thus each As atom loses six 
charges of electricity and each Zn atom gains two, so that it takes 
three Zn atoms to reduce one AsO3 anion. 

(b) In alkaline solution, by means of zinc, aluminium, sodium 
amalgam, or by Devarda’s Alloy (Cu=50, Zn=5, Al=45). This 
reaction also is sometimes attributed to nascent hydrogen: | 


Zn+2Na0H = NaeZnOe+He or Zn+20H  — ZnO2 +He.. 
2Al+2Na0H+2H20 =2N ahOe st Wile or 2Al+20H~ +2Hs0O 
> 2A1097++-3Hpe. 
In the case of Devarda’s Alloy, the reaction is completed mudh’tore 
quickly than by the use of either zine or aluminium alone. Nitrates 


ee ee ee 
: 
“ 


OXIDATION AND REDUCTION 35 


and chlorates may be reduced in a few minutes by means of Devarda’s 


_ Alloy and a few drops of sodium hydroxide; the reaction also takes: 


place in neutral solution, but it takes considerably longer: 
. NO3 +4Zn+70H~ — 4Zn0O2~+2H20+NH3 
ClO3"+3Zn+60H- — 3Zn02*+3H20+CI. 


In the nitrate-ion nitrogen has a positive valence of five; in 
ammonia it has a negative valence of three. By the reduction with 
zinc, therefore, the nitrogen loses eight positive charges, or accepts 
eight electrons. At the same time the zine accepts two electrons, 
forming, in a neutral or alkaline solution, the zincate-anion. Thus 


one atom of nitrogen in the nitrate-ion requires four atoms of zine to 


convert it into ammonia. 

Similarly, the chlorine atom in the chlorate-ion has a positive valence 
of five and is reduced to a negative valence of one by the reaction with 
zine in neutral or alkaline solution. Thus one chlorate-ion reacts with 
three atoms of zinc. Inspection of the above equilibrium expression 
shows that one could predict that the reaction would take place best 
in alkaline solution in accordance with the mass action law. , 

Reduction by means of sulfurous acid takes place in moderately 
acid solution and depends upon the fact that sulfur is more stable 
when it has six positive charges, as in sulfuric acid, than when it has 
only four as in sulfurous acid. Ferric salts are readily reduced by 
this reagent, and since the iron loses only one valence while the sul- 
fur gains two, it is evident that one molecule of sulfur dioxide (the 
anhydride of sulfurous acid) will reduce two atoms of iron in a ferric 
salt: 
Fe2(SO4)3 +2H20+S02 = 2H2804+2FeSOz, 
or 

o¥et * *+-S0."+-Hs0 — 2Fe* *-++-S0.7°+-2H". 


In a similar manner, the arsenate-ion and many other substances 
are reduced very readily by means of SO2 or SO3 ': 


AsOaz +S037>—AsO3_ +5804- 


An excess of aqueous sulfurous acid is added to the solution which is 
to be reduced; it is then heated to boiling; and the boiling is contin- 
ued while a stream of carbonic acid gas is passed through the solu- 
tion until the excess of sulfurous acid is driven off. 

Reduction by means of hydrogen sulfide, in which the sulfur 
atom possesses two negative charges, depends upon its oxidation to 
free sulfur, which is electrically neutral. Thus two atoms of ferric iron 


36 . GENERAL PRINCIPLES 


are reduced to ferrous iron by one molecule of hydrogen sulfide and 

one molecule of potassium dichromate reacts with three molecules of 

hydrogen sulfide: | | 
2Fet*+++HS — 2Fett+2Ht+S 


Cr207-+3H2S+8H* — 2Cr++++3S+7H20. 


One objection to the use of hydrogen sulfide as a reducing agent is 
the difficulty involved in the subsequent removal of the precipitated 
sulfur by filtration. Moreover, hydrogen sulfide is used in qualitative 
analysis chiefly as a precipitant. If a solution contains an oxidizing 
agent (such as nitric acid, chloric acid, chromic acid, etc.), the sulfide- 
ion will be oxidized and there will be separation of sulfur. Any 
_ sulfide obtained will be largely contaminated with sulfur, which renders 
the subsequent examination more difficult. If the solution contains no 
metal which is precipitated by hydrogen sulfide, but contains oxidizing 
agents, it will still cause separation of sulfur. One is often in doubt 
whether there is not some sulfide mixed with the sulfur, and is therefore 
obliged to examine the precipitate further, which is often unnecessary if 
the oxidizing agent is previously destroyed. Hydrogen sulfide reduces 


Halogens: H28+Cle 3 =2HCI+S; 

Nitric Acid: 2HNO3+3H28 =4H,0+2NO0+38; 
Chloric Acid: HC10O3+3H28 =3H20+HCI1+35; 
Ferric Salts: 2FeCls-+ H2S =2HC1+2FeClo-+S; 


Chromic Acid: . 2Cr03+3H2S+6Ht =6H20+2Cr*++t+38; 

Permanganic Acid: 2HMn04+5H28+4H* =8H20+2Mn**+58; 
and many other substances. . 

Reduction with stannous chloride takes place usually in acid 
solutions. The reduction depends upon the fact that stannous ions 
are readily changed to stannic ions: | 

SnCle +Clo = SsnCh, 
or +40 
Snt*++Cle=Snt*t+2Cr, 
Ferric salts, chromates, permanganates, mercuric salts, and many 
others are reduced in this way: 


hed ' 
2Fet*+++S8ntt > 2Fett+Sntt; 
++ 
2CrO4~+3S8n*t ++16H* > 2Crt*+*+3Snt *+8H20; 
+4 
2HgCle +8nt +t — Sntt+t+HgeCle+2Cr; | 


seer iy 
Hg2Cle+Sn** > Sn**+2Hg+2Cr-. 


ee ee ee 

- ae 22 Pat 

a ue Je <<? e 
F 3 2 


ELECTROMOTIVE SERIES AND OXIDATION POTENTIALS 37 


Reduction with hydriodic acid depends upon the change of the 


iodine anion into free iodine. Most substances that are capable of 


being oxidized or reduced readily can be made to react with either 
hydriodic acid or with free iodine. It is easy to detect the presence of 
free iodine, and for this reason the iodometric reactions are extremely 
important in the study of analytical chemistry. To prevent the 
oxidizing effect of free iodine, an excess of potassium iodide is usually 
required, and means are often taken to remove the iodine as fast as it 
is formed; this is in accordance with the mass-action principle, 


2Mn04-+101-+16H* > 2Mn*++8H20+5lo; 
Cr207~+61 +14Ht > 2Cr*+*+*++7H20+3!I2; 
Ket? +1 => Pet t+I. 


Electromotive Series and Oxidation Potentials 


If a substance like sugar lies as a solid on the bottom*of a beaker | 
filled with water, the molecules of sugar tend to distribute themselves 
throughout the solution; in other words the sugar dissolves. The 
tendency of the solid molecules to pass into solution may be regarded 
as the result of pressure and, in fact, it is customary to say that the 
solid substance possesses a solution pressure. 

If sufficient solid is present, eventually, with the aid of diffusion, 
the liquid will reach a state of saturation. The liquid then contains 
an equal quantity of sugar in all its parts and, at the prevailing temper- 
ature, will not dissolve any more sugar. There must, therefore, be 
some force which acts in opposition to the solution pressure and pre- 
vents a saturated ‘solution from dissolving any more of the solid 
substance. This force is the osmotic pressure which the dissolved mole- 
cules exert in the solution. In a saturated solution, the osmotic pres- 
sure, which is itself determined solely by the number of molecules of 
dissolved substance and the temperature, exactly balances the solu- 
tion pressure of the solid substance. The process of dissolving a solid 
substance involves no electrical effects. This is also true when the 
dissolved substance is an ordinary electrolyte, because an equal number 
of positive and negative ions is formed and there is no electric dis- 
turbance. 

The metals themselves, though usually to a much less degree, also 
show a tendency to dissolve when placed in contact with water. In 
this case, however, an oxidation takes place, for, to the extent that it 
dissolves, the metal is converted into electrically chargedions. The 


38 GENERAL PRINCIPLES 


tendency of the metal to dissolve is called its electrolytic solution pres- 
sure. Just as in the case of the sugar, the osmotic pressure of the © 
dissolved ion acts against the solution pressure. The electrolytic 
solution pressure has a definite value which is characteristic of each 
metal. 

' If a metal such as zinc, which oxidizes fairly readily, is placed 
in a saturated solution of zinc sulfate, none of the metal dissolves. If 
it is placed in contact with a dilute solution of zine sulfate, the solu- 
tion pressure of the zine is greater than the deposition pressure and 
some positively charged zine ions pass into solution. Thereby, the 
metal itself acquires a negative charge and the solution a positive 
charge. Asa result of the charge residing upon the zinc ions that have 
gone into solution, an electrostatic force is produced which seeks to 
force the ions back upon the metal. This electromotive force is added 
to the osmotic pressure of all the zinc ions in solution and it increases 
rapidly with the number of ions that dissolve from the metal. When- 
the sum of the osmotic pressure plus the electromotive force is equal 
to the electrolytic solution pressure of the zinc, the zine stops dissolving. 

When a less-readily oxidizable metal, such as copper, is placed in a 
copper sulfate solution the relations are reversed. In this case, except 
in extremely dilute solutions, the osmotic pressure is. greater than 
the solution pressure and the metal does not dissolve; on the contrary, 
a few of the copper ions are discharged on the metal, giving to it a posi- — 
tive charge while the solution becomes negatively charged. Equilib- 
rium is established as soon as a few of the ions have been thus deposited. 
The potential difference between the metal and its solution, or as it is 
often called, the potential of the metal, is said to be positive when the 
charge of the solution is positive; this is the case with the readily oxidiz- 
able metals such as magnesium, aluminium, zinc, iron, etc. On the 
other hand, the potential of the metal is negative when it is difficultly 
oxidizable; this is the case with copper, silver, platinum and gold. — 

Simple contact of a metal with a solution of its ions usually 
results in a potential difference between the metal and the solu- 
tion. Such a potential is determined by the relation that exists 
between the electrolytic solution pressure of the metal and the 
osmotic pressure of the solution. Equilibrium is soon reached in 
most cases and the simple contact of a metal with a solution of its 
ions is not a permanent source of electricity. If, however, two metals 
of different potential are placed in contact with their respective solu- 
tions, then electric charges of different potentials result, and if the two 
metals are connected outside the liquids by a wire, an electric current 
flows from the higher potential to the lower. Since the original differ- 





‘ 


ELECTROMOTIVE SERIES AND OXIDATION POTENTIALS 39 


ences in potential between the solutions and the metals are constantly 
being reéstablished, a permanent current results. This is the prin- 
ciple of the Daniell cell, in which a normal solution of copper sulfate. 
is separated by a porous partition from a normal solution of zine | 
sulfate. A zinc rod is placed in the zinc sulfate solution and a copper 


~ plate in the copper sulfate solution; the current flows through the wire 


from the copper to the zinc and through the solution from the zine 
to the copper. 

Nernst, who was the first to suggest the above explanation of the 
origin of the electromotive force on the basis of the relations of osmotic 
pressure, has worked out a formula for computing the potential differ- 
ence which exists at the place of contact of a metal with a solution of 
its ions. If H denotes this potential in volts, R the gas constant 
expressed in voltsXcoulombs, F the electrochemical equivalent or 
quantity of electricity borne by one equivalent weight in grams of the 
ions of any metal, n the valence of the ions, P the electrolytic solu- 
tion pressure, p the osmotic pressure, and 7 the absolute temperature 
of the solution, the Nernst formula reads: 


eee 
Bey Be 


Substituting the numerical values for R(8.32) and F(96,500), dividing 
by 0.434 in order to use common logarithms, and assuming the ordinary. 
room temperature to be 18° C. (=291° absolute), the formula becomes: 


ata log 2 volts. 





Ege = 


Inasmuch as the osmotic pressure, p, depends solely upon the concen- 
tration of the solution and the temperature, ,the electromotive force 
resulting by the contact of a metal with its ions is shown by the 
formula to increase as the electrolytic solution pressure of the metal 
increases and to decrease asthe concentration of the ions increases. 
Since multiplying a number by 10 simply raises its common logarithm 
one whole unit, the formula also shows that increasing the concentra- 
tion of the solution tenfold lowers the electromotive force in question 
about 0.06 volt when the ions are univalent, 0.03 volt when the ions 
are bivalent, and 0.02 volt when the ions are trivalent. Lowering the 
concentration until it is one-tenth of its original value raises the elec- 
tromotive force nearly 0.06 volt in the case of univalent ions, 0.03 
volt with bivalent ions, and 0.02 volt with trivalent ions. The electro- 
lytic solution pressure is a measure of the readiness with which a metal 
can be converted into its ions; or, since the formation of the ions 


40 GENERAL PRINCIPLES 


involves an oxidation, it determines the readiness with which the ele- 
ment undergoes oxidation. The electromotive force that results can 
be appropriately called the oxidation potential. It is also a measure 
of the force required to deposit a metal from solution by means of the 
electric current. 

It is now easy to understand what happens when a metal is placed 
in a solution containing the ions of some other metal. It is a well- 
known fact that the immersion of a strip of iron in a solution of copper 
sulfate causes the deposition of metallic copper, while an equivalent 
quantity of iron dissolves as ferrous sulfate. The copper is reduced 
to the metallic condition by means of metallic iron and the latter is 
oxidized by means of cupric ions. This is because the electrolytic 
solution pressure of iron is so much greater than that of copper that a 
condition of equilibrium is not reached until practically all of the copper 
has been precipitated. The oxidation potential of metallic iron against 
a molar solution of a ferrous salt is 0.43 volt and of metallic copper 
against a molar solution of cupric ions is about —0.34; the minus sign 
means merely that there is more tendency for copper-ions to be de- 
posited than for metallic copper to pass into solution. The greater the 
positive value of the oxidation potential, the greater the electrolytic 
solution pressure. As the copper is deposited from the solution, its 
oxidation potential becomes gradually larger, and as the iron passes 
into solution its oxidation potential becomes smaller and smaller. The 
Nernst formula shows that the value for metallic copper against a 
tenth-molar solution of its ions will be raised to —0.31 volt and 
the value for iron against a solution of tenth-molar ferrous salt will 
be 0.40 volt. It is evident that equilibrium between the iron and the 
cupric solution will be reached only when the oxidation potential of 
the iron is equal to the oxidation potential of the copper; before this 
happens either the solution will become saturated with ferrous sulfate 
or all but an infinitesimal quantity of the copper will be precipitated. 

It is possible to arrange all the metals in a series according to their 
electrolytic solution pressures. Such a series is called the electro- 
motive series of the metals. It enables one to understand all the 
reactions in which a free metal is involved either as an initial or 
a final product. The entire chemical activity of the metals corre- 
sponds fairly closely with such an arrangement. The members at. the 
top of the series are the most readily oxidizable; those following cop- 
per do not oxidize or rust when exposed to the air. 

The electromotive series shows the relative value of the metals as 
reducing agents. The metals at the top of the series are the best reduc- 
ing agents. Thus the alkalies are such good reducing agents that 





ELECTROMOTIVE SERIES AND OXIDATION POTENTIALS 41 


they will even decompose water at ordinary temperatures, reducing 
the positively charged hydrogen to the neutral condition. 

It is important to remember, however, that it is not alone the electro- 
lytic solution tension which determines the oxidation potential. The 
concentration of the solution also comes into consideration. If the 
oxidation potentials are all measured against equivalent concentra- 
tions, then the order of the metals arranged in the electromotive series 
will correspond exactly to the order of the metals when 


free hydroxyl ions in solution and these are in equi- 
librium with the zinc ions. 


arranged according to the values of their electrolytic ~t®cTRomotive 
: SERIES OF THE 
solution pressures. Mineasay 
‘ We are now able to understand why the alkalies Cashier 
decompose water readily and why the quantity of Rubidium 
zinc ions formed under similar conditions is very Potassium 
small. The oxidation potential of zinc against a molar spgscadl 
~ solution of zinc ions is about 0.76 volt. Water, however, Aluminium 
is ionized very slightly; the table on page 10 gives Manganese 
its ionization as 2X10-’ per cent. The oxidation re 
potential of hydrogen against such a very dilute solu- Thallium 
tion of hydrogen is not 0.0, as given in the table, but Iron 
it is nearer the zinc value. If the ionization of water hae 
were 10-*6, the value would be approximately that of Tin 
zinc. According to the oxidation potentials, therefore, Lead 
we should expect zinc to decompose water with libera- ae 
tion of gaseous hydrogen. As a matter of fact zinc is Bismuth 
oxidized somewhat by contact with water and the Arsenic 
oxidizing agent is the hydrogen of water; but the ee 
reaction does not take place to any extent. The Silver 
primary products of the reactions are zinc ions and Palladium 
free hydrogen, but the escape of the hydrogen leaves Aes 


The table on page 22 gives the solu- 


bility product of zinc hydroxide as 1.8X10-'*. The reason the zine 
does not decompose water, therefore, is because it is protected by 
the film of insoluble oxide or hydroxide which quickly forms upon 
it. On the other hand, when the hydrogen is present in the form of 
an acid, with the anion of which zinc forms a fairly soluble salt, the 
oxidation of the zinc ordinarily continues at the expense of hydrogen 
ions until all of the zinc is dissolved. 

The electromotive series shows the oxidation tendencies of the 
elements, but in attempting to predict what will happen in any given 
case it is necessary to bear in mind that the concentration of the solu- 
tion must be considered and the solubility relations. 


42 _ GENERAL PRINCIPLES 


We have seen that the Daniell cell* is obtained by taking advantage 
of a difference in oxidation potentials, and it was stated on page 29 
that a similar cell could be formed by taking advantage of any reac- 
tion of oxidation and reduction. All reactions of oxidation and reduc- 
tion take place because of differences in oxidation potentials. Just 
as the electromotive series of the metals helps one to predict whether 
a metal will act as a reducing agent or not, so a complete table of oxida- 
tion potentials will help one to determine whether any given reaction 
of oxidation and reduction may be expected to take place in the desired 
direction. Such a table of oxidation potentials is given on page 43. 
The table shows the values referred to molar solutions, the value of 
normal hydrogen ions in the normal hydrogen electrode being taken 
as 0. The positive value of the electromotive force shows that the 
oxidation takes place very readily, the unchanged substance assuming 
a negative charge; or, in other words, if a cell is constructed with 
the normal hydrogen electrode, the direction of the current (positive 
to negative) in the solution is toward the hydrogen electrode when 
the element has a positive potential. The first column in the table 
gives the original state of the element or ion, the second column shows 
the change in charge that this element or ion undergoes, the third column 
the oxidized condition and the fourth column the oxidation potential. 

The electromotive force of any reaction of oxidation and reduction 
is determined by the difference in the oxidation potentials. Thus 
in the Daniell see using molar solutions of copper sulfate and zine 
sulfate, the electromotive force of the entire cell is the difference between 
the oxidation potential of copper (—0.34 volt) and zine (+0.76) =1.10 
volts. 





*Jn the well-known Daniell cell the zinc, which is in contact with zine sulfate 
solution, is the negative electrode, and the copper, which is in contact with 
copper sulfate solution, is the positive electrode; zine passes into solution at one 
electrode and copper is deposited at the other. Outside the cell, the positive-to- 
negative direction is from copper to zinc, but inside the cell it is from zine to 
copper. In measuring the single electrode-potentials of such a cell, all German 
and many English texts assign a negative value to the potential of the element 
which dissolves and a positive value to the element which is deposited. This is, 
however, a purely arbitrary practice, and it seems more natural to assign the 
positive and negative values in the opposite manner. It is rational, in the Daniell 
cell, to assign a positive value to the electromotive force which measures the 
tendency of zinc to form zinc ions, and a negative value to the tendency of cop- 
per to form copper ions. In tracing the course of the current in the Daniell cell, 
it is simplest to start with the zine which forms zinc ions and pass through the 
solution to the copper and back through the wire to the zinc. In the cell, the 
zinc is positive with respect to the copper but in the wire the copper is positive 
to the zinc. 


ELECTROMOTIVE SERIES AND OXIDATION POTENTIALS 43 


OXIDATION POTENTIALS 














Ag Ag 
Original | #8| Higher | E.mf. Original “| 32 | © Higher | Bmt. 
tate ES State. in volts. State. ES State. in volts, 
=| =| 
| SS ae Bet Bava ce +3.03]| Cu®......... a PCr. 52 3: --0.51 
: K° oc, oe AHS 1 Kt rahe eit a a +2 “ 93 Mn0O.+ 3 MnO;- +Ht —0 - 52 
Pea... ¥( Nee Sc. +2.72 20H— - 
Ga aa SV Raw. s: et he, Te § > Ad it Papel gra tao —0.54 
ae VSG OS ie aa +2.7 || Br°+30H- .| 5 | BrO;-+3Ht |—0.60 
oe i 2 |Catt+ +2.6 |) Fett. ..... 1 |Fet+t.....|—0.75 
ee ea a" Po ola etapa y ot en | a. a See a Ot ie ‘1—0.80 
MARE ee 3 | Alt++ 413 oHe® oF 2 |Hett+..... —~0.80 
MPS es. 7 ed 9 a 4d dt BOs oo 2 | O.gas+2H* |—0.80 
Peres oes. VA 07 ae ae +0.76'} Ch° gas+ 2 GHGS: 253 —0.85 
Sa Ae 2 ae ba a ee +0.55 20H 
: a Teale © aman lea 043i Het ii... a are grea —0.86 
2 haa eee Pel oe ey +0.40}| 2Hgt....... 2 | 2Hgt++ —0.92 
Pb°+SO0 = ed ig Pa +0.54)2Br-. oo... 7 Des 3.) ee ae ae ae —1.08 
MUN ci ko 2° 1 Oor®.: +0.23)} Cl-+OH— BS DEON ss StS, —1.10 
ees... Yay is, | a Saree +0.22/| I1°+30H- 5 |10;+8H*.. |—1.19 
Cu°+20H— 2 | Cu(OH), +0.21|} Cré+++ HCrO;-+ {-1.3 
A a Page i sel +0.12 4H.O| 3 4H+ 
Meee iss. . SRE ais. +0.10)} Mnt* + MnO:+ —1.35 
a 3 | Fettt +0.04 2H20}| 2 4H+ 
eee ts.’ DS sieeet atiet +0.00)} 2CI-....... 7 Soe Nh ) PR ee —1.35 
Ct. eek. 1 | Cutt... .}—0.17]| Pht ++ 2 | PbO.+2H+ |—1.44 
Ae°+Cl- ...| 1 | AgCl....... ~ 0.23 20H - 
Hg.Cl, soln. | 2 !2HgCh..... —0.24)) Au®......... j eS sae tes —1.5 
+2Cl— PENS «et Loh, + Vad 61 (2 a ee —1.51 
I°+OH— Oe is eee —0.26)} Mnt++ 5 | MnO, + —1.52 
‘Os Tae 3 | Ast++ —0.29 4H.0 8Ht 
ROU uk. ths > eg 1B ai apa ee —0.34!| MnO.+2H.0} 3 | MnO; + —1.63 
Ag®+2NH;3 1 | Ag(NH;).+ .;—0.38 4H+ 
BERS es AES dar f SOON 2HeO .. 56 oh 2 | H.0.+2Ht .|—1.0 
A eng alas 3 |Cot++. ....|—0.40]| Cott. ..... 1 |Cottt,...|/-1.8 
0) 5 Meas 2 | O.+H,0 —0.41|| Cl, gas+ 4 | 2ClO~ + —1.8 
Sane 3 |Sbtt++t+.....|—0.47 20H— 2Ht 
Br®. liq.+ 5 | BrO;- +H,0|—0.51]| 30,2 gas...... 3 | 203 gas..... —1.9 
30H— See a 2 Fe Ree —1.96 





























The table of oxidation potentials will help to explain many of the 
reactions used in analytical chemistry. All the metals above hydrogen 
will replace the hydrogen of dilute acids; those below hydrogen will 
not do soasarule.. The oxidation of copper to cuprous ions, however, 
corresponds to an oxidation potential of —0.17 volt. When the acid 
is very concentrated and the cuprous solution very dilute, the oxida- 
tion potential of the copper to univalent copper will be above that of 
hydrogen and consequently a little copper should go into solution as 
cuprous salt. In the presence of some oxidizing agent lower than 


44 GENERAL PRINCIPLES 


hydrogen ions in the series, the copper will dissolve readily. Such an 
oxidization agent is the ferric ion, for the table shows that the oxida- 
tion potential of ferrous ions to ferric ions is —0.75 volt; just as cupric 
ions will oxidize metallic iron to the ferrous condition so will ferric 
chloride oxidize metallic copper to the cupric condition: 


Fett *+-+Cu > Cut t+2Fet tf 


The table shows where interference is to be expected. Metallic 
aluminium will precipitate iron from a solution of a ferrous salt. It 
will not precipitate iron if an acid is present because, as the table shows, 
iron will itself liberate hydrogen gas from hydrogen ions; on the other 
hand, from the oxidation potential of ferrous to ferric ions, it is evi- 
dent that metallic aluminium will reduce a ferric solution before 
acting upon the hydrogen ions of the acid. It also shows that metallic 
aluminium will precipitate copper completely even in the presence 
of acid. 

Lead and tin occupy neighboring positions in the series. In a 
neutral solution, therefore, lead will precipitate tin from a solution 
containing stannous ions. As the concentration of the lead ions in- 
creases and the concentration of the stannous ions decreases, the oxida- 
tion potentials approach one another so that equilibrium is soon reached. 
Conversely, when no stannous ions are present, metallic tin will pre- 
cipitate a little lead from a solution containing lead ions, but the equilib- 
rium will soon be reached. ‘The presence of acid, however, will stop — 
both of these reactions. On the other hand, the oxidation potential 
of stannous to stannic tin, although not given in the table, has a nega- 
tive value, and the lead will’reduce stannic ions even in the presence of 
acid. 

All reactions of oxidation and reduction represent reversible reac-: 
tions. The strong reducing agents on being oxidized become weak 
oxidizing agents, and conversely the strong oxidizing agents on being 
reduced become weak reducing agents. Ordinarily hydrogen is 
considered a reducing agent, but when a metal replaces the hydro- 
gen of an acid, hydrogen ions act as the oxidizing agent. 

The mass action law holds for oxidation and reductions as for all 
other chemical reactions that take place in solution; when the oxida- 
tion voltages are far apart, however, the reaction of oxidation and 
reduction will apparently go to completion. 

We shall now turn aside from reactions of oxidation and reduction 
and consider some further applications of the mass-action law as applied 
to reactions that are of importance in analytical chemistry. 





THE EFFECT OF A COMMON ION 45 


The Effect of a Common Ion 


In all equilibrium expressions in which the concentration of an 
ion is involved, the source of the ion is a matter of indifference. Thus 
when sodium chloride and potassium chloride are both present in solu- 
tion, the chlorine ion is common to each. The equilibria between 

the salts and the ions are expressed as follows: 7 
[Nat]x<[Cl]_, [K*] x[CI-] _ 
[NaCl] — tact and [KCl] = Mon 


In each of these equilibrium expressions the concentration of the 
chlorine is the total concentration in the solution. It is evident, 
therefore, that the addition of the extra chlorine ions from potassium 
chloride tends to lessen somewhat the extent of the dissociation of 
the sodium chloride and similarly the presence of the chlorine ions from 
the sodium chloride tends to lessen somewhat the extent of dissociation 
of the potassium chloride. Both sodium and potassium chloride belong 
to the class of strong electrolytes and the dissociation remains con- 
siderable even after the other salt is added. As far as the chemical 
behavior of the salts goes, the effect is scarcely noticeable. As long as the 
solution is reasonably dilute the dissociation of both potassium and so- 
dium chloride will be considerable in spite of the presence of the other salt. 

The relations are quite different when one of the original substances 
is difficultly soluble or only slightly ionized. Thus when a precipitate 
of silver chloride is formed, the solution is saturated with the salt and 
- the solubility product of the ions has been reached. If now a small 
amount of either chlorine or silver ions is added to the saturated solu- 
tion of silver chloride, further precipitation of silver chloride should 
take place. As a general rule, therefore, a slight excess of precipitant 
will make a precipitate less soluble. There are exceptions to this 
rule, however. ‘The rule does not hold if the one of the ions from the 
precipitate shows a tendency to form a soluble complex ion with the 
excess of precipitant. Thus silver chloride is, in fact, less soluble in 
very dilute sodium chloride solution than it is in water, but it dissolves 
in a saturated brine solution more than in pure water, probably owing 
to the formation of a complex ion. Barium sulfate is less soluble in 
dilute sulfuric acid than it is in water, but concentrated sulfuric acid 
dissolves it quite readily. Aluminium hydroxide is precipitated by 
the very careful neutralization of an aluminium salt, but the hydroxide 
dissolves in sodium hydroxide solution, forming sodium aluminate. 
None of these examples is contrary to the mass action law, but they 
show the need of considering all the possible reactions. 

When a weak electrolyte is present in solution, the effect of the 








46 GENERAL PRINCIPLES 


common ion is often quite remarkable. To illustrate this effect let 
us consider the weak electrolytes acetic acid, hydrogen sulphide and 
ammonium hydroxide. The table on page 10 shows that a 0.1 N 
solution of acetic acid is dissociated to between 1 and 2 per cent. The - 
value of the ionization constant at 18° is about 0.000018. From this 
value, the ionization (x) in 0.1 N solution can be computed as follows: 


> =0.000018 - 2=0.00134 = 1.34 per cent. 


The table on page 10 shows that the salts of the type represented by 
sodium acetate are dissociated to about 84 per cent. A liter of 0.1 N 
sodium acetate solution contains, therefore, 0.084 mole of acetate ions, 
whereas one of 0.1 N acetic acid contains only 0.0013 mole of acetate 
ions. When enough solid sodium acetate is added to 0.1 N acetic 
acid to make the solution 0.1 N with respect to both the acid and 
the salt, the common acetate ion tends to repress the ionization of 
both the original molecules; the total concentration of acetate ion 
is increased only very slightly and we are justified in assuming that 
0.084 represents with sufficient accuracy the concentration of the 
acetate ion. The concentration of the hydrogen ions may be des- 
ignated again as x and that of the non-ionized acetic acid as 0.1—z. 
The ionization constant of acetic acid remains the same and the 
mass action expression becomes 


0.084x 
0.1—z 


By adding an equivalent weight of sodium acetate, therefore, the ioniza- 
tion of 0.1 N acetic a¢id is changed from 1.8 per cent to 0.02 per cent. 

The effect of ammonium salt upon the ionization of ammonium 
hydroxide is similar. The table on page 22 gives the solubility 
product of ferric hydroxide as 1.110-%6 and that of magnesium 
hydroxide as 3.410". As the cube of the OH” concentration is 
taken in computing the solubility product of ferric hydroxide and 
only the square of this concentration in the case of magnesium 
hydroxide, the difference in solubilities is not as great as these 
solubility products would indicate, but the values show that it takes 
only an extremely low concentration of hydroxyl ions to satisfy the 
solubility product of ferric hydroxide and many times as much to 
satisfy that of magnesium hydroxide. Ammonium hydroxide added 
to asolution containing ferric and magnesium ions will cause the pre- 
cipitation of both iron and magnesium as hydroxide, but by adding 
sufficient ammonium chloride to the solution, the ionization of the 
base is repressed, exactly in the same way that sodium acetate acted 





= (0.000018; x=0.000021 =0.02 per cent. 





THE EFFECT OF A COMMON ION 47 


upon acetic acid, to such an extent that none of the magnesium 
is precipitated, although the precipitation of the iron remains 
practically complete. 

The effect of hydrogen ions upon the ionization of hydrogen sulfide 
is similar. The solubility product of copper sulfide is 8.510-*5 
and that of zine sulfide is 1.2<10-°%. The precipitation of both cop- 
per and zinc by hydrogen sulfide is practically complete in a solution 
containing no excess hydrogen ions. As the metal precipitates, how- 
ever, hydrogen ions are formed: 


Zn+H.S — ZnS+2H". 


The accumulation of these hydrogen ions serves to repress the ioniza- 
tion of hydrogen sulfide and tends to stop the precipitation of the zine. 
If, therefore, we wish to precipitate copper and leave zinc in solution, 
all that is necessary is to add a little acid at the start; in 0.3 N 
_ acid solution the precipitation of the copper as sulfide is practically 
~ eomplete, while little if any zinc sulfide is precipitated. On the other 
hand, if some sodium acetate is added to the solution, non-ionized 
acetic acid is formed and the accumulation of the hydrogen ions is 
prevented. This effect is so remarkable that it is easier to precipi- 
tate zine sulfide from a solution containing acetic acid and sodium 
acetate than from a solution of zine chloride in water. The concen- 
tration of hydrogen ions is kept very low, even although the solution 
may smell strongly of non-ionized acetic acid. 

The common ion effect is also involved in the solution of precipi- 
tates. The effect is shown, for example, in the solution of calcium 
phosphate. The table on page 10 shows that the tertiary ionization 
of phosphoric acid is about the same as the secondary ionization of 
hydrogen sulfide and the secondary ionization is comparable to the 
primary ionization of hydrogen sulfide. If considerable hydrogen 
ion is added in the form of a mineral acid, the ionization of HPO 
and of HePO4 becomes extremely small. The saturated solution of 
calcium phosphate contains Catt and PO.= ions. These PO4> ions 
must be in equilibrium with added hydrogen ions. When the POs~ 
from HPO,” is kept less than the PO4~ concentration corresponding 
to the value of a saturated solution of calcium phosphate, the calcium 
’ phosphate must tend to dissolve. As Ca(HzPO.)2 is much more sol- 
uble than CaHPOu, which in turn is more soluble than Ca3(PO4)g, it is 
desirable to add enough acid to make the H2POx practically non-ionized. 

The common ion effect which is involved in nearly every reaction 
of precipitation and of solution represents an application of the mass- 
action principle. 


48 GENERAL PRINCIPLES 


Hydrolysis 


Hydrolysis is the name given to the decomposing action of water 
upon many salts. Corresponding to the fact that water is a poor 
conductor of electricity, it follows that water is dissociated only to a 
slight extent. 


H20 = Ht+0OH-. 


According to Kohlrausch and Heidweiler,* the degree of disso- 
ciation at 25° C. is 1.05107; in other words, about 10,000,000 liters 
of water would furnish 1 gm. molecule of ionized water or 1 gm. atom 
of hydrogen ions. Small as this is, it suffices to explain the hydrol- 
ysis of: 

I. The salts of weak acids with strong bases. 
II. The salts of strong acids with weak bases. 

III. The salts of weak acids with weak bases. 

The salts of strong acids with strong bases are not hydrolyzed 
appreciably. 3 

Hydrolysis is shown to take place by the fact that solutions of 
neutral salts corresponding to I react alkaline, those of II react acid, . 
while those of III are sometimes acid and sometimes alkaline. 

The cause of hydrolysis is the action of the ions of water upon 
the ions of the dissolved salt. 

All monobasic salts in aqueous solution are largely dissociated 
into ions (cf. p. 10): 


RA—Rt+A-, 


and the phenomenon of hydrolysis may be represented by the general 
equation 
R*t+A~+H20 s$ ROH+HA. 


ete — +’ 
Salt Base Acid 


In formulating mass-action expressions involving highly ionized 
and slightly ionized substances it has been found best to employ the 
concentration of the ions of highly ionized substances, and the con- 
centrations of the non-donized portions of slightly ionized substances. 
Since in dilute solutions the mass of the water changes inappreciably 
when the reaction takes place, we may neglect it in the formulation. 

There are now three typical’ equations which represent the appli- 
cation of the mass-action principle to the hydrolysis of the above 
three classes of salts. 





* Z. phys. chem., 14, 317. 





HYDROLYSIS 49 


I. Hydrolysis of Salis of Weak Acids and Strong Bases 


The base formed by the hydrolysis is largely ionized, while the 
acid is only slightly so. The mass-action equation now takes the 
form | 
[R*][OH [HA] _[OH7)[HA] 

[R*][A7] 0s aa 

Owing to the presence of the OH ions in appreciable amount, 
all salts of this category react alkaline. The alkali salts of hydro- 
eyanic acid, hypochlorous acid, carbonic acid, boric acid, and hydro- 
gen sulfide are of this type. 





dd) Ki= 


II. Hydrolysis of Salts of Strong Acids and Weak Bases 


Here the conditions are reversed, and it is the acid which is almost 
completely dissociated and the base but slightly. In this case the for- 
mula becomes 


[ROH][H* [A] _[ROH)[A*] 

[R*|[A7] [R*] © 

Salts of this class, such as those of chromium, aluminium, iron, etc., 
react acid when in aqueous solution. 


In the case of polyvalent bases both electrolytic and hydrolytic 
dissociation take place in stages: 





(2) By= 


RCls @ RClot+Cl- 
meme | nog" eenci+c 
| RCIt+2 Rttt+cr. 


RClyt +H.0 @ RClLOH+ Ht 
Hydrolysis RCI* ++2H20 = RCI(OH)2+2Ht 
R**+*+3H20 = R(OH)3+3H". 


Ferric, aluminium and chromic chlorides, for example, react acid 
in aqueous solution. If such solutions be evaporated to dryness, 
considerable hydrochloric acid is volatilized, and the residue obtained 
in an insoluble Haste salt which can only be dissolved by means of 
acid. 


III. Hydrolysis of Salts of Weak Acids and Weak Bases 


The acids and bases formed by hydrolysis are only slightly dis- 
sociated, but to different extents. Our formula becomes 


[ROH][HA] 
[R*][A7] 





(3) Ky= 


50 GENERAL PRINCIPLES 


Salts of this type are especially subject to hydrolysis and, as both 
the acid and base are slightly ionized, the hydrolysis may take place 
to a considerable extent without the soluten manifesting either acid 
or basic properties. 

If the electrolytic dissociation of the acid is greater than that 
of the base, the solution reacts acid; and conversely, when the base 
is stronger than the acid the solution of the salt shows an alkaline 
reaction. 

Neutral ferric acetate in a boiling, aqueous solution is hydro- 
lyzed: 


Fe(C2H302)3+2H20 @ Fe(OH)2C2eH302+2HC2Hs02. 


Basic ferric acetate is precipitated and can be removed by filtering 
the hot solution. If the solution is allowed to cool, the reaction tends 
to take place in the reverse direction and some of the basic salt goes 
into solution. Heat and dilution always favor hydrolysis.* 


IV. Hydrolysis of Salts of Strong Acids with Strong Bases 


Salts of this type yield, by hydrolysis, acids and bases which are 
almost entirely dissociated in dilute aqueous solution, and the general 
equation becomes: 


[R* OH |[H*[A*] 
[R*][Aq] 


The H and OH ions, however, are in equilibrium with undissociated 
water; the solution reacts neutral and contains only as many H and 
OH ions as correspond to the ionization of water, which is so small 
that there remains only the electrolytic dissociation of the salt to — 
be considered. 





Ky = =[H*][OH"]. 


Salts of this type are not subject to appreciable hydrolysis. 


The hydrolytic action of water, as well as the mass-action law, 
may be illustrated by the following experiment: A little water, added 
to a solution of antimony chloride in concentrated hydrochloric acid, 
causes precipitation of antimony oxy-chloride: 


—Cl 
Sb—Cl+}j, 02 = Sb_ oy +2HCI, 


—Cl 





*The ionization constant for water is 0.0:;12 at 25° but rises to 0.015 at 100°. 
The water is, therefore, much more dissociated at the higher temperature and 
as a result the hydrolysis is favored. 








HYDROLYSIS 51 


which dissolves on adding a little concentrated hydrochloric acid. 


Further addition of water again precipitates the basic salt, which will 
dissolve in more of the concentrated acid, etc. It is obvious that by 
increasing the mass action of the water the reaction goes from left to 


_ right, while by increasing the concentration of the hydrochloric acid 
it goes from right to left. 


The analytical chemist frequently desires to assist or to prevent 


hydrolysis. The mass-action principle shows how this can be done. 


To assist a chemical reaction it is necessary to increase the concen- 
tration of one of the original reacting substances or to lessen the con- 
centration of one of the substances formed. Aside from the reac- 
tions of oxidation and reduction, all the reactions that take place 
completely in aqueous solution are those in which the concentration of 
one of the substances formed is practically negligible. This is due to 
the formation of (1) a precipitate, (2) a gas, or (3) an undissociated sub- 
stance (cf. p. 26). Neutralization takes place betweenanacid and a base 
because of the tendency to form undissociated water. Hydrolysis is 
the reverse of neutralization and is due to the fact that water is slightly 
dissociated. Hydrolysis takes place when one of the products is a 
gas, a precipitate, or an undissociated substance. Thus the hydrolysis 
of a salt of a weak acid and a strong base is due to the fact that the 
dissociation of the weak acid is slight. The hydrolysis of a salt of 
a strong acid and a weak base is due to the formation of the undisso- 
ciated base. Hydrolysis takes place most readily when both the acid 
and the base are weak, because then both the H ions and the OH ions 
are removed from the solution to form undissociated acid and undisso- 
ciated base. Hydrolysis of a salt of a strong acid and a strong base 
cannot take place because there is then no tendency for the H and OH 
ions of water to be removed. ‘To assist hydrolysis, boiling is advisable, 


- because the water is so much more dissociated at this temperature 


than when cold. The reaction that takes place on boiling often 
proceeds in the other direction on cooling, simply on account of the 
change in the ionization of the water itself. Dilution favors hydrolysis 
because it diminishes the concentration of the substances formed; the 
concentration of the ions from water is not changed, but the relative pro- 
portions of these ions to other ions present is increased. Hydrolysis is 
also favored when one of the products is gaseous or insoluble; thus 
ammonium carbonate is very easily hydrolyzed, because both ammonia 
and carbon dioxide are volatile, and ferric acetate is hydrolyzed not 
only because the acid and base are both weak, but the basic ferric 
acetate is also very insoluble. Finally, to stop hydrolysis it is only 
necessary to add a little free acid or free base at the start, and the 


52 GENERAL PRINCIPLES 


efficiency of the acid or base is proportional to the extent to which 
it is itself ionized. If, the solution becomes alkaline as a result 
of hydrolysis, then a little alkali will best stop hydrolysis, but if the 
solution becomes acid, a little acid should be added. Moreover, it is 
advisable to work in cold and concentrated solutions. A dilute © 
aqueous solution of potassium cyanide has the odor of free hydro- 
cyanic acid and reacts alkaline to litmus, but if a little caustic potash — 
is added, the hydrolysis of the salt is prevented and the odor of hydro- 
cyanic acid can be detected no longer. Similarly, boiling a solution 
of neutral ferric sulfate results in the precipitation of a basic salt, but 
the precipitation does not take place if a little sulfuric acid is added 
at the start. | 


Amphoteric Electrolytes 


According to ideas that prevailed nearly a century ago, an acid was 
regarded as composed of a negative element and hydroxyl; a base was 
regarded as composed of a positive element and hydroxyl. According 
to the modern conception, an acid is a substance capable of yielding 
hydrogen ions and a base is a substance capable of yielding hydroxyl 
ions. The so-called negative elements are sometimes simple negatively- 
charged ions, as the Cl” of hydrochloric acid, and sometimes complex 
anions, such as the SO," of sulfuric acid, containing a positively charged 
element and negatively charged oxygen. 

The extent to which the ionization takes place in the solution | 
of an acid or a base varies greatly. In 0.1 N solution, the table 
on page 10 shows that some acids and some bases are dissociated to 
about 90 per cent, while other acids and other bases are dissociated 
to 1 per cent or less. We are accustomed to regard the strength of 
an acid or of a base as shown by the extent to which ionization takes 
place in aqueous solution. 

As a general rule, when the positive charge on an element is in- 
creased, the more difficult it becomes for the corresponding hydroxide 
to ionize as a base and the greater the tendency for the hydroxide © 
to ionize as an acid. Thus the higher the positive charge on the 
atom, the more negative it becomes according to the old-fashioned 
conception. This apparent contradiction is not quite as inconsistent 
as it seems; it arises from a confusion of the intensity factor of 
electric energy with the quantity factor. All forms of energy are 
composed of these two factors and such confusion of the factors is quite 
common. 

The quantity of electricity corresponding to a unit electric charge 
on a gram atom of any element is 96,500 coulombs. All univalent 


AMPHOTERIC ELECTROLYTES Ne BS 


a elements bear this charge and all bivalent elements twice as much. 


_ On the other hand, the tendency of the atoms to accept or give up 


electrons varies greatly, as the table of oxidation potentials shows 
 (p. 48). It is customary to speak of the elements which show the 


| greater tendency to give up one or more electrons of negative electricity 
| as more positive than those which show less tendency to lose electrons. 
_ In other words, the elements nearer the top of the electromotive series 
are regarded as more positive than those which are below them in this 





series. ‘Thus the older conception of positive and negative nature 
was largely one of electromotive force, or voltage, rather than one of 
electric nature or quantity of electricity. It corresponds to the 
behavior of the element in an electric couple such as that of the Daniell 
cell; the zinc becomes the anode in such a cell and the copper the 
cathode; the current flows through the cell from the zine to the cop- 
per and it is customary to say that the zinc is positive to the copper. 
The student should be careful to distinguish in his own mind the dif- 
ference between the meaning of the word positive when used in this 
sense of voltage and when used to represent the character of an elec- 
tric charge. 

Water is a substance which ionizes to a slight extent and the initial 
products of the ionization are H+ and OH-. Water, therefore, may 
be considered as being both an acid and a base. There are other 
substances which act as acids without forming at one time any more 
hydrogen ions than does water and other basic substances which do 
not form more hydroxyl ions. Water is characterized by the fact 
that it forms an equal amount of both hydrogen and hydroxyl ions 
by its primary ionization. Other hydroxides are known which have 
both acid and basic properties, but this is due to two distinct kinds of 
ionization. At one time they dissociate as an acid and at another 
time as a base. Such substances are said to be amphoteric electrolytes. 

Aluminium hydroxide is an amphoteric electrolyte. This sub- 
stance has a very small solubility product and the quantity of ions 
present in the saturated solution’ is extremely small. The chemical 
behavior of aluminium hydroxide shows; however, that it is capable 
of dissociating in two ways: 


Al(OH)3 — Al(OH)2*-+OH™ (ionization as a base); 
Al(OH)3 — Ht+AlO3H2- — Ht+AlO2~+He20 (ionization as an acid). 


The mass action principle enables one to predict which of these 
ionization reactions will take place. In the presence of a strong acid, 
such as hydrochloric acid, the common ion effect of the hydrogen 


/ 


54 GENERAL PRINCIPLES 


ions will prevent the ionization of the aluminium hydroxide as an 
acid; the effect is much more marked than that of sodium acetate 
upon acetic acid, for in this case the difference in the percentage 
ionization of hydrochloric acid and aluminic acid is much greater. 
On the other hand, the presence of the hydrogen ions favors the ion- ~ 
ization of the aluminium hydroxide as a base. To establish the proper 
equilibrium between Ht, OH~ and H20, nearly all the OH™ ions from 
the aluminium hydroxide react with the H* of the hydrochloric acid. 
On account of this removal of the OH~, the ionization of the Al(OH)s 
progresses, the Al(OH)2* ionizes into Al(OH)t* and OH™ and finally 
the Al(OH)** ionizes into Al**+ and OH. In this way Al(OH)s3 
dissolves to form AICls. 

Similarly, in the presence of sodium hydroxide, the common ion 
effect. of the OH™ prevents the aluminium hydroxide from ionizing 
as a base and the H* ions formed by its dissociation as an acid react 
with the OH™ ions from the sodium hydroxide, and the final result is 
that the aluminium all dissolves as sodium aluminate, NaAIOzg. 

Whenever an oxide or hydroxide dissolves by chemical reaction 
with an acid and also by chemical reaction with a base, the substance 
is obviously an amphoteric electrolyte. 


Detection of Acids and Bases. Theory of Indicators * 


To detect the presence of free hydrogen cations or of hydroxyl 
anions, certain colored, organic substances are used, called indicators. 
These indicators are very weak acids, very weak bases or amphoteric 
electrolytes and the free acid, or base, is an unstable substance 
which tends to undergo a slight rearrangement of the atoms in order 
to assume a condition of greater stability. The color of any organic 
compound is due to a certain special arrangement of certain atoms, 
the chromophor; when this arrangement is changed, the color is also 
changed or lost. 

Methyl orange is an amphoter and is capable of forming salts 
with both acids and bases, but its indicator characteristics are due 
to its very weak basic properties. The neutral solution of its sodium 
salt is used as an indicator. In this sensitive neutral solution we have 
a condition of equilibrium between the two isomeric forms of methyl 
orange as expressed by the equation 


HSO3 E CgHsN :N- CeH4N(CHs)2 = OsteBane -N: CeH4 : NC 








*Cf. J. Srizaciitz, J. Am. Chem. Soc., 25, 1117; 39; Acres, tbid., 37, 39, 42; 
McCoy, ibid., 31, 508; Saum, Z. phys. Chem., 57, 471; and A. A. Noyzrs, J. Am. 
Chem. Soc., 32, 815. 








DETECTION OF ACIDS AND BASES 55 


_ The compound on the left is yellow in color and its color is due to 
the azo group N : N; the other compound is red, having for its chromo- 
phor the quinoid group : CgHa:. 
The sodium salt of methyl orange is yellow and has the formula 
NaSO3 : CeHaN :N- CeH4N(CHs)z2, 
and when decomposed by acids the free sulphonate 
tril on gaa : CeH4 ; N(CHs3)2 
| 





is formed, which is red. 

The red quinoid form is ionized as a weak base and forms red salts 
with acids. It does not form salts readily with weak acids, such as 
carbonic or acetic acid, because, as we have seen, salts of weak bases 
and weak acids are hydrolyzed. This is why methyl orange is not a 
sensitive indicator for weak acids. As a very weak base it will be driven 
readily out of its red salts by other bases, even weak ones; and the free 
base will revert again to its yellow form, the result being that methyl 
orange is an excellent indicator for weak bases. 

Phenolphthalein, another valuable indicator, is a very weak 
acid. The free acid, however, is unstable, and when set free from 
one of its colored salts reverts instantly into a colorless lactoid form: 


HOOCCsHsC(CeH.OH) . Caki«:s Me See Cer as 
Re 


Colorless 





The red color is in this case also due to the quinoid grouping: CgHa:. 
In the free acid, the condition of equilibrium favors the lactoid form, 
and only minimal quantities of the quinoid acid are present. This 
trace of quinoid acid is ionized and in equilibrium with its ions: 


HOOCC.H4C(CgH40H) : CeHs : OR 
Ht+O0CC.gH1C(CgH:0H) > CeHs: OO. 


The addition of an alkali causes the hydrogen ions to disappear, 
more of the quinoid molecules must be ionized to preserve equilib- 
rium, and the quinoid molecules in turn be reproduced from the lactoid 
as fast as the former are converted into the salt. Phenolphthalein is 
@ very sensitive indicator toward acid, but on account of being such 
a weak acid it does not form stable salts with weak bases. 

Besides these indicators, others are often employed, among which 
may be mentioned Litmus and Lacmoid, which are red with acids 
and blue with alkalies; and Turmeric, which is brown with alkalies 
and yellow with acids. 


‘ 


56 GENERAL PRINCIPLES 


The various indicators, therefore, differ from one another with 
regard to the extent to which they normally undergo ionization either 
as an acid or as a base. They show, as a result, specific degrees of 
sensitiveness to hydrogen cations and to hydroxyl anions. The follow- 
ing table shows the concentration of the ion required to produce the 
color change. The tables refer to results obtained when about two 
drops of a 0.1 per cent solution of the indicator is added to 10 ce. of | 
the solution tested. 


SENSITIVENESS OF INDICATORS * 











Concentration Required 
for Color Change of 
Indicator. Color with H+. Color with OH-. 
Ht, OH-. 
Phenolphthalein.....| colorless........ pink (3 PA, 0.071 0.041 
Azolitmin (in litmus).| violet pink...... WIGOLS «hae s 0.051 0.051 
Methyl orange...... reddish-orange..| yellow....... 0.0.1 0.01 

















The table shows that of these three indicators phenolphthalein is 
the most sensitive to acids, and methyl orange is the most sensitive to 
bases. | 


The Evaporation of Acids 


It is often necessary to diminish the acid concentration of a solu- 
tion. The simplest way to accomplish this is by neutralization, but 
it is often undesirable to introduce foreign substance into the solution, 
and in such cases the acid is removed by evaporation. To prevent loss 
by spattering, it is well to evaporate in a porcelain dish on the steam 
bath; but, to save time, the evaporation may take place over a free 
flame. In this case the solution should be kept in motion either by 
stirring or by rotating the contents of the dish and moving it back and 
forth over the flame. For evaporations over a free flame, an Erlen- 
meyer flask is often used; the’sides of the flask, being cooled by the 
air, act as a condenser and delay evaporation, but they also serve to 
prevent mechanical loss. Spattering is caused by steam being formed 
too rapidly at the bottom of the dish, by steam being formed beneath 
a crust which has formed on the surface of the liquid, or by steam being 
formed from enclosed mother-liquid in a solid that has separated. The 
steam escapes then only after its pressure has become considerable. 
The breaking of glass dishes when heated over a free flame is due to 
overheating the glass by contact of the flame with some part that is 





*From J. Stieglitz: Qualitative Analysis. 


THE EVAPORATION OF ACIDS | 57 


not being kept relatively cool by contact with liquid. Moving the 
vessel back and forth over the flame prevents breakage by preventing 
overheating of the glass; the bottom is kept uniformly wet on the inside 
even when but little liquid remains. 

The behavior of acids upon evaporation is interesting. A solution 
of hydrochloric acid of specific gravity 1.10-boils at 110° under atmos- 
pheric pressure. The solution contains 20.2 per cent of hydrogen 
chloride and its concentration is thus a little less than 6 N. No 
other mixture of water and hydrochloric acid boils as high and con- 
sequently in evaporating a solution of hydrochloric acid, although the 
total quantity of hydrochloric acid constantly diminishes, the concen- 
tration of the remaining acid will tend to approach the composition of 
the constant boiling mixture. If the acid is more dilute at the start, 
'_ evaporation will cause the acid that remains to be more concentrated, 
and if more concentrated at the start, evaporation will cause the 

remaining acid to be more dilute. . 

Nitric acid similarly forms a constant boiling mixture with water. 
This mixture boils at 120.5° and contains 68 per cent of nitric acid; 
the acid concentration of the mixture is about 15.5 N. 

Sulfuric acid forms with water a constant boiling mixture that con- 
tains 98.3 per cent of anhydrous sulfuric acid and boils at 330°. It was 
once thought that these constant boiling mixtures of acid and water 
corresponded to definite hydrates of the acid, but the composition of 

such a mixture and the boiling-point varies with changes in pressure in a 
way that proves no definite compound is present. 

When it is desired to change a solution of a chloride to one of a 
nitrate it is very easy to do this by evaporating once or twice with 
nitric acid. ‘The nitric acid reacts with the hydrochloric acid, 








6HCI+2HNO3 =4H20+2N0+3Clh, 


* » and by adding an excess of the nitric acid all of the chloride is decom- 
: posed. The most economical way to accomplish the change is to 
evaporate nearly to dryness and then add a little strong nitric acid, 
repeating the evaporation and addition of acid until no more red fumes 

are evolved on adding the acid (NO-+air-red NOz). 

Similarly a solution of nitric acid may be changed to one of hydro- 
chloric acid. Aqua regia is formed, as in the above case, and by using 
an excess of hydrochloric acid all the nitrate is decomposed. 

The change of a solution of nitric acid or of hydrochloric acid to 
one of sulfuric acid is based on another principle. In this case the 

change is accomplished by evaporating till dense fumes of sulfuric acid 


58 GENERAL PRINCIPLES 


are evolved. Very little sulfuric acid is lost until all of the more volatile 
acid has evaporated. | 

Phosphoric acid is still less volatile. By evaporation with phos- 
phoric acid even sulfuric acid can be distilled off without losing an 
appreciable quantity of phosphoric acid, but the phosphoric acid is 
changed to pyrophosphoric acid or to metaphosphoric acid. 

To change a sulfate or a phosphate solution to a chloride or nitrate 
presents a more difficult problem than that of simple evaporation. It 
is customary in such cases to remove the phosphate or sulfate ions by 
precipitation or to precipitate the desired substance and to dissolve it 
in the desired acid. 


Filtration and Washing of Precipitates 


When a precipitate is produced in qualitative analysis it is necessary 
to remove it from the liquid by means of filtration. The size of the filter 
used should be determined by the size of precipitate and not by the 
volume of liquid. In attempting to detect the presence of traces of a 
substance it is often necessary to work with large quantities of the orig- 
inal substance, and this involves the use of correspondingly large 
volumes of liquid. If from such a solution a few milligrams of a pre- 
cipitate is formed, it will be practically lost if spread over a large filter, 
and further work with the precipitate is hampered. 

Before examining such a precipitate it must be completely freed 
from all traces of the filtrate. This is accomplished by washing. Wash- 
ing must be continued until no test can be obtained with the wash- 
water for a certain substance known to be present in the filtrate. For 
example, suppose it is necessary to filter off some suspended barium sul- 
fate from a solution containing sodium sulfate; the precipitate must be 
continuously washed until a sample of the wash-water, acidified with 
hydrochloric acid, no longer gives a precipitate on the addition of barium 
chloride. As a rule, it is not advisable to run the wash-water into the 
filtrate, because this occasions an unnecessary dilution of the latter. 
The filter must always be smaller than the funnel, and the precipitate 
should not extend higher than to within 5 mm. of the top of the filter. 
Finally, large precipitates should be avoided as far as possible, for they 
render exact work more difficult—filtration and washing consuming 
too much time. 


Colloidal Solutions 


If an alkaline silicate in dilute solution is mixed with dilute hydro- 
chloric acid, the solution remains clear and apparently unchanged in 





hE a a 


COLLOIDAL SOLUTIONS 59 


spite of the fact that silicic acid is only slightly soluble in water. If the 
liquid thus obtained is placed in a vessel whose walls are formed of 
parchment paper or of bladder and this vessel is placed in pure water, 
it is possible to wash out all the excess hydrochloric acid and all of the 
sodium in the form of sodium chloride. This is the so-called process 
of dialysis. The liquid in the dialyzer contains silicic acid. 

‘The silicic acid, however, is not in the form of a true solution. If it 
is evaporated, an amorphous, gelatinous mass is obtained which is only 
slightly soluble in water. The boiling-point and freezing-point of the 
solution do not differ much from that of pure water. The addition of 
various substances, especially of salts, causes the liquid to solidify as a 
jelly, especially if it has been concentrated somewhat by evaporation 
in the cold. 

A similar experiment can be performed with arsenious acid and 
hydrogen sulfide. The aqueous solution of arsenious oxide turns 
orange yellow when hydrogen sulfide is passed into it and becomes 
opalescent without forming any precipitate, although arsenious sulfide 


is only very slightly soluble in water. The liquid will pass unchanged 


through an ordinary paper filter. If a little hydrochloric acid or some 
salt solution is added to the yellow opalescent liquid, a heavy precipi- 
tate of arsenious sulfide at once forms. 

A liquid in which a very insoluble substance appears to be in solu- 
tion far beyond its usual degree of solubility, and yet does not show at ~ 
all the behavior of an ordinary supersaturated solution, is said to con- 
tain the substance in colloidal solution. 

Such solutions are very commonly formed with difficultly soluble 
gelatinous substances. It was formerly thought that such solutions 
represented true solutions of difficultly soluble substances in the form 
of a soluble colloidal modification, but more recent observations with the 
ultramicroscope have indicated that colloidal solutions are really sus- 
pensions of minute solid particles. The colloidal condition evidently 
interferes with the precipitation of insoluble substances, and since most 
of the separations of analytical chemistry depend on the formation of 
precipitates, analytical chemistry is chiefly concerned with the colloidal 
condition as one that is to be avoided as much as possible. 

The suspended particles of most colloids carry electrical charges; 
a potential difference exists between the particles and the liquid. Some- 
times the charge on the particles is positive, as with the colloidal solution 
of ferric hydroxide, aluminium hydroxide and chromium hydroxide; 
and sometimes the charge is negative, as with the colloidal solu- 
tions of silicic acid, stannic acid, the sulfides of arsenic and cadmium, 
silver iodide, silver chloride, gold, platinum and silver. Substances 


60 GENERAL PRINCIPLES 


in the colloidal condition which carry an electric charge are precipitated 
by an electrolyte. Negatively-charged colloids are precipitated by the 
action of positive ions and, conversely, the positively-charged colloids 
are precipitated by negative ions. The precipitated substance carries 
with it a little of the precipitating ion in the form of an adsorption com-— 
pound. The precipitating power of electrolytes increases decidedly 
with the valence of the precipitating ion. 

Some colloids do not carry much of an electric charge and sometimes 
the charge is positive in acid solutions and negative in alkalies. Col- 
loids which do not carry much electric charge are not precipitated 
by dilute solutions of electrolytes. These colloids are coagulated by 
heat, by any other method of dehydration such as the addition of con-— 
centrated salt solutions, or by the addition of another solvent such as 
alcohol. Colloids which are not sensitive to precipitation by electro- 
lytes often act as protective colloids and hinder the precipitation of 
other substances. Thus tannic acid, gelatin and albumin act as such 
protecting agents, probably by forming protective films. 

To prevent the colloidal condition in analytical work it is necessary 
to take certain precautions. *Since electrolytes cause the precipitation 
of colloids, it may be expected that their absence will tend to aid in 
the formation of colloidal solutions. Thus in washing a precipitate 
which is known to have a tendency to pass into colloidal solution, it is 
best to wash with a solution of some indifferent electrolyte rather than 
with pure water. Again, if precipitations are attempted in very dilute 
solutions of such substances as arsenious sulfide, it is best to have an 
electrolyte in the solution at the start. It is also necessary to remember 
that the colloids carry down with them the precipitating ion by which 
they are coagulated. To avoid analytical mistakes caused by such loss 
of ions, a sufficient concentration of some other ion is usually provided 
in the form of an acid or an ammonium salt. The washing of the 
precipitated colloid with ammonium chloride or ammonium nitrate 
gradually removes these precipitated ions, but it is very difficult to avoid 
all danger of loss from this source. 

When protective colloids, especially of the gelatin or albumin type 
are present, they may interfere so much with the common precipitation 
tests that it is necessary to destroy the protective colloid, usually by 
oxidation, before proceeding with the analysis. 








. 
v 


REACTIONS IN THE DRY WAY 61 


II. REACTIONS IN THE DRY WAY 


These reactions are employed chiefly in the so-called ‘‘ preliminary 
examination,” in testing the purity of precipitates, and in the examina- 
tion of minerals. The most important reactions of this nature consist 
in the testing of a substance with regard to its— 

1. Fusibility; 

2. Ability to color the non-luminous Bunsen flame; 
3. Volatility; 

4. Behavior toward oxidation and reduction. 

In order to carry out these reactions it is customary to use the 
non-luminous gas flame; and to understand the operations to be 
described it is necessary for us to know something about the composi- 
tion of illuminating gas and the nature of the flame. 

The illuminating gas of Zurich averages the following composition: 
CO2=2.0 per cent; C,H2,=4.5 per cent; O2=0.2 per cent; CO=8.0 
per cent; H2=48.0 per cent; CH4=33.0 per cent; and No=4.3 per 
cent. 

All these components, except CO2, O2, and Ne (which are present 
only in small amounts), are combustible; they are reducing substances. 
illuminating gas ordinarily burns with a luminous flame, and the 
luminosity is due to the presence of unsaturated hydrocarbons (C,Ho,), 
principally ethylene, propylene, acetylene, benzene, etc. If ethylene 
is heated to a certain temperature, it is decomposed into methane and 
carbon: 

C2oH4=CH4+C, 


and it is glowing carbon which causes the luminosity of the flame. | 

The other unsaturated hydrocarbons behave like ethylene. The 
remaining combustible constituents of illuminating gas burn with a 
non-luminous flame. If we bring air into the gas, the flame becomes 
non-luminous. With the Bunsen burner air is introduced by opening 
the holes at the base of the burner. In such a gas-flame there are, 
according to Bunsen, the following parts (Fig. 1*): 

I. The inner cone of the flame, aab, in which no combustion takes 
place, because the temperature here is too low. This part of the flame 
contains unburned gas mixed with about 62 per cent of air. 

II. The flame mantle, indicated by acaba, which is composed of 
burning gas and air. 





*In the drawing, the lines d represent a metallic chimney, or flame-protector, 
which rest upon the support ee. It is advisable to furnish each burner with such 
a flame protector. 


62 GENERAL PRINCIPLES 


Ill. The luminous tip, at b, which does not appear unless the air- 
holes are closed somewhat. 

In these three principal parts of the flame Bunsen distinguished 
Six reaction zones: | 

1. The base of the flame at a’. The temperature here is relatively 
low, because the burning gas is cooled by the constant current of 
fresh air, and also because the burner itself conducts away considerable 
heat. This part of the flame serves to test volatile substances to see 
whether they impart color to the 
flame. In case several substances are 
present which color the flame, it is 
often possible to observe the colors 
one after the other, in that the most 
volatile substance colors the flame 
first, and later the colors caused by 
the less volatile ones are seen. This 
would not be possible at a hotter part 
of the flame, as all of the substances 
would then be immediately volatilized, 
producing a mixture of colors. 

2. The fusing zone at B. This lies 
at a distance of somewhat more than 
one-third of the height of the flame, 
and equidistant from the outside 
and the inside of the mantle, which 
is broadest at this part. As this is 
the hottest part of the flame (about 
2300° C.), it serves for testing sub- 
stances as to their fusibility and. 
i volatility. 

Fic. 1. 3. The lower oxidizing flame lies in 

the outer border of the fusing zone at 

y, and is especially suited for the oxidation of substances dissolved 
in vitreous fluxes. 

4. The upper oxidizing zone, at ¢, consists of the non-luminous 
tip of the flame, and acts strongest when the air-holes of the lamp 
are fully open. It is used for various oxidizing tests, the roasting 
away of volatile products of oxidation, and generally for all processes 
of oxidation where the very highest temperature is not required. 

5. The lower reducing zone lies at 6, in the inner border of the fusing 
zone next to the dark cone. As the reducing gases are mixed here with 
oxygen from the air, many substances which are reduced by the upper 


©) 


























REACTIONS IN THE DRY WAY 63 


reducing flame are unaffected in this zone. This part of the flame is 
consequently very well adapted for a test which cannot be made 
with the blowpipe, namely, reduction on the charcoal stick, and in 
vitreous fluxes. 

6. The upper reducing flame is at ny, in the luminous tip of the dark 
inner cone, which may be produced by gradually diminishing the supply 
of air. If the luminous tip has been made too large, a test-tube or 
porcelain dish filled with water and placed over it will be blackened,” 
which should never be the case. This luminous tip contains no free 
oxygen, is rich in separated incandescent carbon, and has, therefore, 
a much stronger reducing action than the lower reducing zone. It is 
used more particularly for the reduction of oxides collected in the form 
of incrustations. 


METHODS FOR THE EXAMINATION OF A SUBSTANCE IN THE Dry Way 


1. TEestT OF THE FUSIBILITY 


This test is principally made in the examination of minerals, which 
are introduced into the flame in the loop of a platinum wire (about as 
thick as a horsehair). The sample is examined, after heating, by means 
of a magnifying glass to see whether the corners are rounded, due to 
melting. The potentially hottest temperature of the fusing zone 
amounts to about 2300° C.* It will never be possible to reach this 
temperature with the test, because the substance itself loses heat by 
radiation. As the amount of heat lost by radiation is proportional 
to the surface exposed, it is evident that we will obtain the maximum 
heat by using a very small sample and holder. For this reason a coarse 
wire should not be used for this test. | 

We distinguish the following degrees of heat: 


1. Faint red glow.......... 525°C. Melting-point of magnesium.. 631° 
2: Dark red glow:..:....... 700° aluminium 658.7°T 
3. Bright red glow.......... 950° silver..... 960.5° 
TRO: 1063° 
Moe GOW PIOW So eee ds 1100° Soca.  1083° 
5. Faint white glow........ 1300° nickel.... 1452° 
6. Full white glow.......... 1500° platinum., 1755° 





* This temperature will be considerably lower with too large a supply of air. 
According to Naumann, the temperature of illuminating gas with 1} times its 
volume of air reaches about 1818° C., but the temperature obtained is usually 
lower owing to loss by radiation. The finest platinum wire can be melted by 
means of the flame, but not when it is as thick as a horsehair. 

+ Circular No. 35 of the Bureau of Standards, Washington, D.C. Cf. Burgess- 
IeChatelier: “ High-Temperature Measurements.” 


64 GENERAL PRINCIPLES 


Below 525° C. the following substances melt: tin at 232°, bismuth — 
at 271°, lead at 327°, zine at 420°. 


2. CoLoR IMPARTED TO THE FLAME 


The substance (best in the form of the chloride) is placed in the 
loop of a fine platinum wire, introduced into the base of the flame, 
and then finally brought into the fusing zone. 


3. TEST OF THE VOLATILITY 


A little of the substance is heated in a small test-tube or in a piece of 
glass tubing sealed at one end. Volatile substances are transformed 
into vapor, often without melting, and the vapors condense on the 
cooler walls of the tube. 


4. OXIDATION AND REDUCTIONS 


(a) In Vitreous Fluxes or Beads 


To make a bead, borax (Na2B407:10H2O) or salt of phosphorus 
(NaNHs4HPO4,+4H20) is used. A piece of very fine platinum wire, 
about 3 em. long, is sealed into the end of a glass tube. The wire is 
heated to redness, and then quickly dipped into the borax or salt of 
phosphorus, held near the flame, whereby a small amount of the 
salt isfused to the end of the wire. By repeated heating and dip- 
ping into the salt a bead of sufficient size is obtained. This should be 
about 1.5 mm. in diameter at the most.. It is not advisable to make a 
loop at the end of the wire, because in this way the exposed surface is 
unnecessarily increased. There is no danger of the bead falling off, 
provided the wire is held horizontally in the flame and the bead is not 
too large. In order to bring the substance in question into the bead, 
it is only necessary to moisten the latter with the tongue, and then dip 
it into the finely powdered substance, which will cause a small amount 
to adhere to the bead. It is preferable to introduce too little substance 
into the bead rather than too much, because, in the latter case, the bead 
will become dark and opaque. The oxidation of the substance in the 
bead is brought about by heating it in the lower oxidizing flame; reduc- 
tion is usually effected by heating in the lower reducing zone, and cool- 
ing in the dark inner cone, in order to prevent oxidation, which might 
take place if the substance were cooled in the air. 

In order to clean the wire, a borax bead is produced on the wire, 
which is then heated, as is shown in Fig. 2, a, on one side of the bead 


s 





- REACTIONS IN THE DRY WAY 65 


only, so that the latter runs along the wire in the opposite direction, 
dissolving off all impurities. By heating the bead from the other 
side, Fig. 2, 6, it is driven toward the end of the wire, from which it 
can be shaken off by a quick jerk. By repeating this process three 
times the wire is cleaned with the exception of a small amount of 
adhering borax-glass, which can be removed by heating the wire in 
the fusing zone until the sodium flame entirely disappears. 




















Fig. 2. 


(b) Reduction on the Charcoal Stick 


These exceedingly beautiful reactions are among the most sensi- 
tive of those used in analytical. chemistry, and should be faithfully 
practiced by every beginner. The cause of their sensitiveness is ' 
due to their taking place on the extreme end of a tiny piece of char- 
coal, that is at a point, so that the sample has no opportunity to spread 
itself over a large surface, which is the case with the ordinary reactions 
on charcoal before the blowpipe. 

To carry out these reactions, we use an ordinary splinter consisting of 
good, straight fibers, such as used in an old-fashioned brimstone match 
(not a safety-match, which has already been subject to chemical treat- 
ment). It isimpregnated with sodium carbonate (soda) in the following 
manner: A crystal of sodium carbonate (NazCOs-10H2O) is warmed 
in the flame, whereby it melts in a part of its water of crystallization. 
Three-fourths of the length of the match is now smeared with this liquid 
soda, and the match is then slowly rotated on its axis in the flame, 
until the soda:melts and penetrates the charcoal. On withdrawal 
- from the flame there should be no place which continues to glow; 
should the latter be the case, the stick should be quickly immersed in 


66 GENERAL PRINCIPLES 


the soda again. In this way one obtains a solid little piece of charcoal, 
which can be heated for a long time without burning through. 

In order to carry out a reduction, a small amount of the substance 
to be examined is mixed on the palm of the hand with an equal amount 
of calcined soda, a small drop of melted soda is added, and the mixture 
is made into a paste by means of the blade of a penknife. The warmed — 
piece of charcoal is then rubbed into the mixture, which adheres to it: 
The sample is first heated in the lower oxidizing flame until it has melted, 
and then moved into the lower reducing flame. The reduction will 
be made evident by a violent swelling up of the melt, caused by the 
evolution of carbon dioxide. As soon as the mass melts quietly the 
reduction is complete. The substance is allowed to cool in the dark 
cone, after which it is removed from the flame. The metal is now found 
on the extreme end of the carbonized match, concentrated in a point. 
This point is broken off, and triturated with a small amount of water 
in an agate mortar. The excess of sodium carbonate goes into solution, 
part of the charcoal floats on the surface of the water, while the heavier 
metal sinks to the bottom. In case the reduced metal is iron, nickel, 
or cobalt, it will not be noticeable to the eye, but it may be taken up 
with a magnetized knife-blade, to which it will adhere, usually mixed 
with charcoal. This should be dried by cautious warming, the tuft 
of metal taken off, rubbed between the thumb and forefinger, and 
then brought into contact with the knife again, to which only the metal 
will now adhere. The metal is then transferred to a piece of washed- 
filter-paper about 3-4 mm. wide and 50 mm. long, so that it comes as 
near as possible to the end of the strip. By means of a capillary tube, 
a drop of hydrochloric acid and one of nitric acid are added, and the 
paper is warmed over the flame until the black speck (the metal) has 
disappeared, when the final test can be made. 

In order to test for iron, a drop of potassium ferrocyanide is added, 
whereby the presence of iron is shown by the appearance of a distinct 
formation of Prussian blue. To test for nickel and cobalt, the metal is 
dissolved in nitric acid, the excess of acid is evaporated off, and a drop 
of concentrated hydrochloric acid added, whereby the paper is colored 
blue if cobalt is present; the nickel shows at the most only a very weak 
greenish color—usually, however, no color. A little caustic soda 
solution is now added, and the paper held in the vapors of bromine; 
in case either nickel or cobalt is present a brownish-black spot appears, 
due to the formation of either Ni(OH)3 or Co(OH)s. 

If, however, the metal reduced was malleable, it is usually obtained 
in the form of a metallic globule on the end of the match, where it can 
be examined with the aid of a lens. Copper is not always obtained as a 


i 
o 
4 
; 





REACTIONS IN THE DRY WAY 67 


globule, but usually as a reddish, sintered mass. By pressing down on 
a malleable metal in the agate mortar it is obtained as a glistening 
fragment, which can be readily separated from the specifically lighter 
charcoal by washing. To accomplish this the agate mortar is inclined 
and a stream of water is directed sideways upon the mass, whereby 
the charcoal is washed out with the water, and the metal is left clean. 
It is transferred to a watch-glass and tested as follows: 

1. The Metal is White (Pb, Sn, Ag, Pt). The metal is treated 
with a few drops of nitric acid and carefully warmed. Lead and silver 
dissolve readily, particularly upon addition of a little water. Silver 
will be detected by the addition of a drop of hydrochloric acid, whereby 
white silver chloride, soluble in ammonia, is precipitated. The test 
for lead is dilute sulfuric acid, which precipitates white lead sulfate. 

If the metal, on treatment with nitric acid, remains unchanged, 
it is probably platinum. It should be dissolved in aqua regia, evapor- 
ated to dryness, dissolved in a little water, and potassium chloride solu- 
tion added. A yellow, crystalline precipitate confirms the presence of 
platinum. Ifthe metal, when treated with nitric acid, becomes changed 
into a white insoluble oxide, it is tin. In this case, another fragment of 
metal is dissolved in concentrated hydrochloric acid and tested for tin 
by means of mercuric chloride solution, or by a solution of bismuth 
oxide in caustic soda. 

2. The Metal is Yellow to Red (Cu, Au). Copper is readily dissolved 
in nitric acid, and the solution gives with potassium ferrocyanide a 


reddish-brown precipitate. Gold is insoluble in nitric acid, but soluble 


in aqua regia. The evaporated solution gives a violet-brown color with 
stannous chloride, due to finely divided gold. 


(c) Reduction in a Glass Tube 


Besides the borax bead and the charcoal stick, reduction is often 
effected by means of metallic sodium, potassium, or magnesium. 

Thus small amounts of phosphorus in anhydrous salts may be 
detected in the following manner: The substance to be tested is placed 
in a glass tube, 3 mm. wide and 50 mm. long, which is closed at one end. 
A small cylinder of potassium or sodium (freed from petroleum by 
rubbing between filter-paper), or even a piece of magnesium wire, is 
added to the tube, and the contents then heated until the glass itself 
begins te soften. The reaction is so violent that the substance seems to 
take fire. After cooling, the tube is broken in a porcelain mortar, 
when by breathing over the mass, the smell of phosphoretted hydrogen 
may be detected. 


68 GENERAL PRINCIPLES A, 


The halogens, sulfur, and nitrogen are tested for in a similar 
way, as will be shown later. 


(d) Reduction in the Upper Reducing Flame for the Purpose of Forming 
Metallic and Oxide Incrustations 


The volatile elements which are reducible by means of hydrogen 
or charcoal may be detected in this part of the flame with the great- 
est ease, as, for example, arsenic, anti- 
mony, cadmium, bismuth, selenium, and 
tellurium. The metallic incrustations are 
obtained by holding in one hand a small 
portion of the substance on a thin asbestos 
thread (platinum will be attacked) in the — 
upper reducing zone of a small gas flame, 
where the oxide is reduced to volatile 
metal, and burned in the upper oxidizing 
flame to oxide. In the other hand, closely 
over the substance to be tested, is held 
a glazed porcelain evaporating-dish, filled 
with water, as is indicated in Fig. 3 at B. 
The metallic vapors are condensed by 

Fic, 3. the cold dish, and deposited on it in 

the form of a metallic mirror or film. 

If, however, the dish is held above the upper ‘oxidizing flame 

(at A), there is formed a thin, often invisible, oxide incrustation on 
the bottom. 

Should it be necessary to treat the metallic incrustation with a 
large amount of solvent (as is necessary in the detection of selenium 
and tellurium), the porcelain dish is replaced by a test-tube half filled 
with cold water. A somewhat larger test-tube is used to hold the sol- — 
vent, and the smaller test-tube, on which the incrustation was deposited, 
is placed within the larger tube and the liquid warmed if necessary. 




















(e) Blowpipe Reduction on Charcoal 


These tests are made in the so-called “ preliminary examination.” 
For this purpose a small cavity is made with a penknife in a piece 
of good charcoal (preferably of linden wood), in which a knife-bladeful of 
the substance to be tested is placed, previously mixed with twice as much 
anhydrous sodium carbonate. As charcoal is a porous substance, it will 
readily absorb melted substances, such as salt of the alkalies. Other 





REACTIONS IN THE DRY WAY 69 


substances are changed, by means of the sodium carbonate used, into 
carbonates, which are, for the most part, decomposed, on heating, into 
oxides and carbon dioxide. The oxides of the noble metals are decom- 
posed, without the aid of the charcoal, into oxygen and metal; while 
those of the remaining metals are either reduced to metal or remain 
unchanged. Thus CuO, PbO, BizOz, Sb203, SnOe, Fe203, NiO, and 
CoO are reduced either to a fused metallic globule (Pb, Bi, Sb, Sn, Ag, 
and Au), or to a sintered mass of metal (Cu), or to a glistening metallic 
fragment (Fe, Ni, Co, Pt). The oxides of zinc, cadmium, and arsenic 
do not give metallic globules, but are, however, easily reduced to metal. 
These metals are so volatile that they are changed into vapors, and are 
carried from the reducing zone of the flame into the oxidizing zone, where 
they are changed into difficultly volatile oxides. These oxides, which 
have characteristic colors, are then deposited on the charcoal outside 
the cavity. 

Zine gives an incrustation which is yellow while hot, and white 
when cold; that of cadmium is brown; while the oxide of arsenic gives 
a white and readily volatile incrustation. Furthermore, the volatiliza- 
tion of arsenic gives rise to a characteristic garlic odor. The metals 
lead, bismuth, and tin give, besides the metallic globule, an oxide 
incrustation which is typical. 

At the same time, nitrates, nitrites, chlorates, etc., may be recognized 
by the fact that they cause a very rapid combustion of the glowing — 
charcoal (deflagration). This deflagration is not to be confused with 
a decrepitation which takes place on heating substances containing 
enclosed moisture or gases, such as rock salt, fluor-spar, etc. Crystals 
of such substances are burst by the quick eee of the enclosed 
liquid, and scattered about. 

Many difficultly fusible substances do not melt into the charcoal. 
Thus many silicates form a bead with the soda, which only after con- 
tinuous heating will give up the alkali and allow it to be absorbed by 
the charcoal, leaving behind the white infusible silica. Phosphates 
and borates act similarly, only these do not leave behind an oxide, 
but a fused glass. Infusible white oxides, as those of calcium, stron- 
tium, magnesium, aluminium, and many of’the rare earths (Welsbach 
mantle, for example), glow very brightly, and in fact more brightly 
as they are more strongly heated. 


70 GENERAL PRINCIPLES 


Division of the Metals into Groups 


The metals, for purposes of analytical chemistry, may be divided 
into five groups: 

The First Group contains those metals whose chlorides are insoluble, 
or difficultly soluble, and whose sulfides are insoluble in dilute acids. 
‘They may, therefore, be precipitated from their solutions by means 
of either hydrochloric acid or hydrogen sulfide. 

The Second Group contains those metals whose chlorides are soluble, 
but whose sulfides are insoluble in dilute acids. They may be precipi- 
tated from their solutions by means of hydrogen sulfide, but not by 
hydrochloric acid. 

The Third Group contains those metals whose sulfides are solu- 
ble in dilute acids, but are insoluble in water and alkalies; and-also 
those metals whose sulfides are hydrolytically decoupanet into 
hydrogen sulfide and insoluble hydroxide. The members of this group 
are precipitated completely by hydrogen sulfide only from alkaline 
solutions. 

The Fourth Geos contains those metals whose sulfides are soluble 
in water, but whose carbonates are insoluble in the presence of ammonium 
chloride. They are precipitated by ammonium carbonate in the presence 
of ammonium chloride, but not by any of the above reagents. 

The Fifth Growp contains magnesium and the alkalies; they are 
not precipitated by any of the above reagents. 





In order to carry out an analysis with certainty it is necessary 
to understand not only the reactions of the different elements, but 
we must know as well the sensitiveness of each reaction. The analyst 
should be able to draw a conclusion by the size of the precipitate formed 
as to the approximate amount which is present in the original substance. 
This, however, is possible only when the experiments are made with 
known amounts. Consequently reagents of a known strength are used 
and allowed to act on known amounts of the different substances. Ac- 
cording to the suggestion of R. Blochmann* it is well to make the solu-: 
tions of the different reagents either double-normal, normal, half- 
normal, or tenth-normal. For many years the author has used in his 
laboratory solutions of reagents and salts according to this principle, 


* Berichte, 1890, 31. 





THE LABORATORY REAGENTS 71 


and has found that the beginner in this way gets a far better under- 
standing of the stoichiometrical relations than when solutions of almost 
any concentration are used, as was formerly the custom. 

By a normal solution is understood one which contains in a 
liter one gram-equivalent of the substance in question, referred to a 
gram atom of hydrogen as a unit. A tenth-normal solution will contain 
one-tenth of a gram-equivalent in a liter, ete. 

- Thus one liter of a normal solution will contain 

















HCl = 36.46 gms. 
H2SO4 im 98.08 _ 49 o4 ems. 
2 2 
Ere 08-08 55’ 6g gms. 
3 3 equivalent to one 
NaOH = 40.06 gms. gram-atom of hydrogen. 
KMn0s _ 158.11 _ 5) 69 gms. 
5 5 
Bor _ AE = 49.08 gms. 


The great advantage of this system is that one always knows how 
much of one solution should be used in order to react with another 
quantitatively. Thus 1 cc. of a normal caustic soda solution will 
neutralize 1 cc. of any normal acid, or 2 cc. of any half-normal acid. 
In the same way 1 cc. of a normal solution of sulfuric acid, or of any 
sulfate, will precipitate quantitatively the barium from 1 cc. of a normal 
barium chloride solution. 


The Laboratory Reagents 


I. CoNcENTRATED ACIDS 








Sp.gr. Per cent by Wt.| Approx. Conc. 
MRMINOCTUOEIC BON oe dcocp.b 06 5 aces bess 1.19 37.9 12 N. 
MMA LIT SO ONO 06 109 5 a PE AO ue dis een ern eee’ vat, Tt cate etnintets 24.N. 
MN IMR N Shas ashe bo slnk ee FR acy ales o's 1.42 69.8 16 N. 
MMIIOTIO BOWE oc.) bse ke we ss eles ee © 0 hoe 85 15.N. 
Sulfuric acid........ eats ey OLY 1S Re gee 1.84 96.0 36 N. 














* This is considering the solution as an oxidizing agent, As a precipitant, 
the normal solution of KMnQ, would contain one mole per liter. The reagent 
is almost inyariably used as an oxidizing agent and not as a precipitant. 



























































72 GENERAL PRINCIPLES 
II. Diturep Acips 
Sp.gr. Per cent by Wt.} Approx. Conc. — 
POCO BOG... |: . sce has Scalise 1.04 34.6 6N. 
Hydrochloric acid 5.00 80... ee eee 1.10 20.0 6N. 
PRISE OIE oe. Ca Es aka «bd Galas alee 1.20 32.3 6 N. 
PePGHOriG Hid © 0. oe ee PS 1.12 tinea ‘2N. 
SEO MOME Ooi 5+. dies beets ato lee aes 1.18 24.8 6N. 
PEUPIIUONIS MONG oy. fs Soy, oie eis nd (Satd. soln. of SO2) 0.33 N. 
SRPERPIO MIDI Se wicca oe wee (150 g. per liter) 2N.° 
III. Basss 
Sp.gr Per cent by Wt. Approx. Conc. 
Ammonium hydroxide................ 0.90 28% NHs 15 N Ys 
Ammonium hydroxide...........:.... 0.96 9.9%NHs 6 N. 
Barium hydroxide. .................:. (satd. soln. Ba(OH)2-8H2O) 0.4 N. 
Potassium hydroxide.................. 1.37 36.9 9 N. 
EIN DVOTORIGC as ses cae 1,22 19.7 6 N. 
IV. Sauts © 
(a) Ammonium Salts 
Name of Salt. Directions for Making Reagent. Cone 
Ammonium acetate Mix equal volumes of 6-normal acetic acid 3.N. 
NH,C2H;02 or NH,Ac and 6-normal ammonia. 
Ammonium carbonate | Dissolve250gms.freshly powderedammonium 6N. 
(NH,)2COs carbonate in 1 liter 6-normal NH,OH. The 
commercial salt is NH4sHCO;3+NH,CO.N Ha. 
Ammonium chloride | Dissolve 54 g. in 1 liter of water. N. 
NH.Cl 
Ammonium molybdate | Dissolve 75 g. of pure ammonium molybdate} 0.85 N. 
(NH,)2MoO, in 500 cc. water, pour the solution into 500 
cc. of 6-normal- HNO; and shake till the 
precipitate dissolves. “The formula of the 
commercial salt is (NH4)sMo7O.4-4H20. 
Ammonium oxalate | Dissolve 35 g. (NH4)2C20,-H2O in 1000 cc. | 0.5 N. 
(NH4)2C204 water. 
Ammonium polysulfide | Digest 1 liter 6-normal ammonium monosul- 6N. 
(NH,4)2Sz fide with 25 gms. flowers of sulfur for some 
hours and filter. 
Ammonium sulfide Pass H.S into 200 ec. 15-normal NH,OH in 6N. 
(NH,).S a bottle immersed in running water or in 
ice water until no more gas is absorbed; 
then add 200 cc. 15-normal NH,OH and 
dilute to 1 liter. ; 





























THE LABORATORY REAGENTS 73 
(b) Other Salts 
Name of Salt. Formula. Wooue Nest ead Cone 

EE RSE Nhs role aa» ao KAI(SO4)2-12H20...... 475 59 0.5 N. 
Barium chloride........ BaCh:-2H.0.....5....0. 244 122 N. 
Bismuth nitrate........ Bi(NOs3)3-5H2O0........ 484 80 0.5 N. 
Cadmium nitrate....... Cd(NO3)2°4H2O....... 308 154 N. 
Cadmium sulfate....... 38CdS0O,4-8H20......... 770 64 0.5 .N. 
Calcium chloride....... CaCl, -6H20........... 219 110 N. 
Calcium sulfate........ AMM IAM LEM. bere s 2 oboe 172 2.6 Satd. 
Chrome-alum.......... KCr(SO,4)2-12H2O...... 399 83 0.5 N 
Cobalt nitrate........:. Co(NOs3)2:6H20........ 291 50 1% Co 
Copper sulfate......... CuSO,-5H20.......2.. 250 125 N. 
Ferric chloride......... POC GH i odes ss 270 90 i a 
Lead acetate........... Pb(C2H302)2-3H20..... 379 190 N. 
Magnesium sulfate. .... MgSO,-7H20.......... 247 123 N. 
Manganese sulfate...... MnS0O,-4H20.......... 223 56 0.5 N. 
Mercuric chloride....... POCA as gaia se Sand 8 ow ah 272 27 0.2 N. 
Mercurous nitrate.... ~. Hg2(NOs)e.......--.-. §25 263 N. 
Nickel sulfate.......... NiSO,-7H20.......... 281 70 0.5 N. 

, Potassium bichromate.. .| K;Cr.O7.............. 294 49 N.f 

- Potassium chromate... .| K,CrO,............... 194 291 3.N.? 
Potassium cyanide...... FECA ee ats on Din aus ys 65 65 N. 
Potassium ferricyanide. .| K3Fe(CN)............. 329 110 N.* 
Potassium ferrocyanide..| K,Fe(CN).s-3H2O...... 422 105 N.* 
Potassium iodide....... 1 10 Ma ee ere es geen ae mk heeaien 166 17 0.1 N. 
Potassium nitrite... .... BEV ee St kare ee elak 85 255 3 N. 
Potassium permanganate} KMnO,.:............. 158 10 0.3 N.f 
Potassium thiocyanate. .| KCNS................ 97 97 aS. 
Silver nitrate.......... WISIN erectus peleacate b 170 17 0.1 N. 
Sodium acetate......... NaC2H;02-3H20....... 136 136 N. 
Sodium arsenite........ PRG ics oes ehe ek 130 130 N. 
Sodium bromide........ NaBr-2H,0 or NaBr.. ./139 or 103} 69 or 51 | O.5 N. 
Sodium carbonate...... INN ac heie nals Dees 7 106 159 3 N. 
Sodium hypochlorite....| NaClO............... 74 37 N.f 
Sodium nitrite......... INGINONS Sais Fate Weenie} «8 2 69 207 3N. 
Sodium phosphate...... NaeHPO,:12H20....... 358 119 N. 
Sodium thiosulfate...... NaeS.0;-5H20......... 248 124 0.5 N.f 
Uranyl acetate. .4..../1 UO2(C2H302)2+2H20. ake 425 21 0.1 N. 





* As precipitant. 


Pt) in 100 ce. water. 


t As oxidizing agent. 


ft As reducing agent. 


V. SpeciAL REAGENTS 
Chloroplatinic acid: Dissolve 26.5 gms. H2PtCl,-6H2O (corresponding to 10 gms. 


Kther saturated with HCl: Saturate anhydrous ether at 0° with dry HCl gas. 
Ferrous sulfate, 2 N: Dissolve 280 gms. FeSO,-7H2O in 6 N oa acid _and keep 
in contact with iron nails. 


Hydrogen peroxide, 3 per cent. 


fre 


ee 


User RA 


Magnesium ammonium chloride, N in MgCl: ‘Dissolve 100 gms. MgCl, -6H2O and 
100 gms. NH,Cl in water, er, add 50 cc. 15-normal NH,OH and dilute to one liter... 

Potassium mercuric iodide, 0.5 normal in K,HglI,: Dissolve 115 gms. HglIy and 80 
gms. KI in enough water to make the volume 500 cc.; add 500 ce. 6-normal 
NaOH and decant the solution from any precipitate that may form on standing. 


used, as well as the solids, must be absolutely free from ammonia. 


The stock solution should be kept in the dark. For sensitive work the water 


74 GENERAL PRINCIPLES 


Potassium pyroantimonate: Add 20 gms. of the best commercial salt to 1 liter of 
boiling water, boil until nearly all the salt has dissolved, cool quickly, add 30 
ec. of 10 per cent KOH solution, and filter. : 

Sodium cobaltinitrite: Dissolve 250 gms. NaNO: in 500 ce. water, add 150 ce. 
6-normal HC,H;O2 and 25 gms. Co(NO;3).-6H2O. Let the mixture stand over 
night, filter and dilute to 1 liter. 

Stannous chloride, l-normal: Dissolve 113 gms. SnCl.-2H:O in 100 cc. 12-normal 
HCl, dilute to 1 liter and keep in bottles containing a strip of pure tin. 

Starch and potassium iodide: Rub 20 gms. soluble starch to a thin paste with. 
a little water in a mortar and pour the paste into 1 liter of boiling water. 
Boil five minutes and filter through a loose plug of cotton wool. Add 10 gms. 
KI and 5 ce. chloroform to the filtrate. ~ 

Turmeric: Shake turmeric powder with 95 per cent alcohol and filter. 

Urea: Dissolve 200 gms. urea in 1 liter 6-normal HCl. 


VII. SatTuraTeD SOLUTIONS 


Barium hydroxide ................ 
Bromine water). F6 Ase 
SO oe UE eh eee a 
Chiorme water: . 60044 S.icc 
Hydrogen sulfide.................. 
BAVC WRU i sk bos coo we Gtr ees 


'a\e + oe © bie # 


1000 cc. water dissolves. 
68 gms. Ba(OH)2 at 20°. 
32.68 gms. Bre. 

2.6 gms. CaSO,-2H,0. 
6.5 gms. Cle. 


VII. Sprcrau SoLvENtTS 


Amy] alcohol. 
Carbon disulfide. 
Chloroform. 


Ethyl alcohol (95 per cent). 
Methyl alcohol (free from acetone). 


VII. Souip REAGENTS 


Absorbent cotton 

Ammonium chloride, NH,Cl 
Ammonium nitrate, NH,NOs3 
Barium carbonate, BaCOs 

Borax, NazB,O; 

Boric acid, HBO; 

Calcium chloride (anhydrous) 
Calcium oxide 

Calcium phosphate 

Chloride of lime, CaOCl-Cl 
Copper wire or turnings 

Ferrous sulfate, FeSO, -7H,O 
Ferrous sulfide, FeS : 
Iron nails, Fe 

Lead (finely granulated) Pb 

Lead dioxide (free from Mn), PbO: 
Litmus paper, blue 

Litmus paper, red 

Paraffin 

Potassium acid sulfate (fused), KHSO, 
Potassium carbonate, K,COs 


Potassium chlorate, KCIO; 

Potassium dichromate, K2Cr,O,7 

Potassium ferricyanide, K;Fe(CN)g 

Potassium iodide, KI 

Potassium nitrate, KNO; 

Silica (precipitated), SiOz 

Silver sulfate, AgeSO, 

Sodium acetate, NaC.H;0, 

Sodium ammonium phosphate (micro 
cosmic salt), NaNH,HPO,-4H,O 

Sodium bismuthate, NaBiO; 

Sodium carbonate, NazCOs 

Sodium peroxide, Na,Oz2 

Sodium nitroprussiate, 
Nazk’e(CN)s ° NO -2H.O 

Sodium sulfide, NaS -9H:O 

Sodium tetraborate, (see Borax) 

Starch, (CeHiOs)~ 

Tartaric acid, H2CsH,O¢ 

Tin (finely granulated) 

Zine (finely granulated) 


DETERMINATION OF THE SENSITIVENESS OF REACTIONS 75 


Determination of the Sensitiveness of Reactions. 


| The more sensitive a reaction is, the smaller will be the amount 
_ of the substance which can be detected in a given volume, in a definite 
time, with the reagent in question. Let us assume that the amount 
of substance taken is dissolved in 100 cc. of liquid, the time allowed: 
- to be two or three minutes, and the limit of sensitiveness to be the 
smallest amount of substance which can be detected under these con- 
ditions. 

A few examples will make the method clear: 

Magnesium salts are precipitated by means of sodium phosphate, 
in the presence of ammonium chloride and ammonia, in the form of 
magnesium ammonium phosphate. What is the sensitiveness of this 
reaction? We take 1 cc. of our normal magnesium sulfate solution, 
add three drops of ammonium chloride solution, and two or three 
drops of ammonia and sodium phosphate solutions; the characteristic 
white precipitate is formed immediately. We dilute, now, the normal 
solution of magnesium sulphate ten times, and repeat the experiment 
with 1 cc. of the diluted solution. The result will be— 





1 ce. of N. Mg solution, 100 cc. =1.2 g. Mg, reacts immediately. 

mee. .-N/10;: Mg “ 100 ec. =0.12 “Mg, 

icc. N/100 Mg e 100. cc: =0:012. -- “* Mg, ~. * s 

lec... N/1000 Mg 100 cc.=0.0012 “ Mg, ‘*  afterafewseconds. 

lec. N/10000 Mg - 100 cc. =0.00012 ‘ Mg, “ after one or two 
minutes. 


If, therefore, 100 cc. of a solution contain 0.00012 gm. Mg, the 
magnesium can be detected within one or two minutes. Should the 
detection of smaller amounts be desired, the solution must be con- 
centrated by evaporation. 

This reaction can be called very sensitive. The following potassium 
reactions are much less delicate: 


(a) Reaction with Chloroplatinic Acid (page 78) 


1 cc. of 0.2 N KCl solution, 100 cc. =0.78 gm. K, reacts with a drop 
of H2PtCle immediately. 

1 cc. of 0.02 N KCl solution, 100 cc.=0.078 gm. K, does not cause 
precipitation within three minutes. . 

1 cc. of 0.04 N KCl solution, 100 cc. =0.156 gm. K, does ee cause 
precipitation within three minutes; but does, however, on addition 
of two drops of alcohol. 

1 ce. of 0.06 N KCl solution, 100 cc.=0.234 gm. K, reacts imme- 
‘diately on stirring. 


76 GENERAL PRINCIPLES 


The sensitiveness of the reaction lies, therefore, between 0.156 and 
0.234 gm. K per 100 cc. In order to detect smaller amounts of potas- 
sium than 0.156-0.234 gm. per 100 cc., the solution must be strongly 
concentrated by evaporation 


(b) Reaction with Tartaric Acid (page 79) 


1 cc. of 0.2 N KCl solution, 100 cc.=0.78 gm. K, reacts immedi- 
ately with two drops of sodium acetate and two drops of a concentrated ° 
solution of tartaric acid. 

1 cc. of 0.02 N KCI solution, 100 cc.=0.078 gm. K, reacts after 


one to two minutes with vigorous shaking. 

This can be taken as the limit of sensitiveness. 

If the beginner will test the delicacy of reactions in this way, he 
will quickly get a clear insight into the solubility relations of the 
different salts. 

Probably the quickest way of learning these relations is to analyze 
first of all solutions known to contain an equal quantity of each constit- 
uent of a given group. It is well for the beginner to start with 100. 
ec. of solution and 0.1 gram of each dissolved cation or anion. 

When the sensitiveness of the reactions is known, qualitative 
tests are often the most accurate methods for estimating small quantities 
. of substances. Thus it is possible to detect very small quantities 
of titanium by means of the hydrogen peroxide test, and the most — 
accurate method for determining such small quantities is to compare 
the intensity of the color change with that produced in a series of — 
solutions containing known quantities of titanium. 








7 


PART II. REACTIONS OF THE METALS (CATIONS) 


The separation of the metals into groups (cf. p. 70) is based upon 
the varying solubilities of the chlorides, sulfides, hydroxides and car- 
bonates. In general, the metals removed first in the scheme of analysis 
are those which form the smallest number of soluble salts and the metals 
tested for last are those which form the largest number of soluble salts. 
In qualitative analysis, the relatively insoluble salts occupy the chief 
interest and it is, therefore, simplest to begin the study of the subject 


with that group of metals which is tested for last. Compounds con- 


taining these metals are very common and are used to some extent in 
the separation and identification of the metals of the other groups. A 
knowledge of the characteristic reactions of such compounds is neces- 
sary in order to understand the chemistry involved in the analysis of 
the other groups, and this furnishes another important reason for 
taking up the study in the reverse order to that in which the analysis 
is usually carried out. In this part of thé book only the reactions of 
the more common elements will be considered. The rarer elements, 
such as lithium, cesium and rubidium of the alkali group, will be con- 
sidered in Part V. 


GROUP V. THE ALKALI GROUP 


POTASSIUM, K; SODIUM, Na; and AMMONIUM, NH4 


The metals potassium and sodium are the most reactive of all the 
common positive elements (cf. p. 40). They oxidize very rapidly when 
exposed to the air and decompose water at ordinary temperatures; the 
hydrogen of water is reduced to the gaseous condition and the metal 
is oxidized to alkali hydroxide in equilibrium with alkali cations and 
hydroxyl anions. On account of the extent of the ionization (cf. p. 10) 
the alkali hydroxides form very strong bases. The solid hydroxides 
are the most stable of all hydroxides; they do not break down into oxide 


and water even on being melted. The pure oxides are difficult to pre- 


pare; cautious heating of the metals in air results in the formation of ' 
considerable peroxide as well as oxide. 
77 


78 REACTIONS OF THE METALS 


Ammonium is classed with the alkali metals because the solubility © 
of ammonium salts is similar to that of potassium salts. The ammonium 
radical differs from the alkali metals in being capable of oxidation, the 
usual product of the oxidation being nitrogen gas. Ammonium hydrox- 
ide is a much weaker base than the other hydroxides, and the salts are 
either volatile or are decomposed on being heated. The salts of potas- 
sium, sodium and ammonium are colorless for the most part, and 
readily soluble in water. Of these salts the carbonates, the tertiary 
and secondary phosphates, the cyanides, and the borates react alkaline in — 
aqueous solution (hydrolysis). The salts of the alkalies are more or 
less volatile and impart to the non-luminous flame characteristic colors. 
When a solid is involved, either as initial substance or as final product, 
in any of the characteristic reactions of this and the following groups, 
it will be designated usually by bold-faced type in the equation. The 
formation of a gas will usually be indicated by placing an arrow after 
the symbol in an equation. 


POTASSIUM, K. At. Wt. 39.10 
Sp. Gr. 0.87. M. Pt. (Melting-point) 62.5° C. 


Occurrence.—Sylvite (KCl), isometric, and carnallite (MgCl, 
KCl-6H2O) orthorhombic, occur at Stassfurt in the presence of halite 
and anhydrite. Saltpetre (KNOs), orthorhombic prisms. Further, 
in very many silicates, e.g., monoclinic feldspar (KAISisOg), and 
muscovite (KHe2AlsSigOi2); also in plants in the form of organic 
salts, which yield on combustion potassium carbonate (potash). 


REACTIONS IN THE WET WAY 


| Potassium forms very few salts that are difficultly soluble in water. 
The chloroplatinate, acid tartrate, and perchlorate are the least soluble, 
and are consequently used in the detection of potassium. 
1. Chloroplatinic Acid,* Hy2{PtCl¢|], gives in concentrated solutions 
of the chloride a yellow precipitate of potassium chloroplatinate, 


PtCh-"+2Kt—KoPtCly, 


which consists of small regular octahedra (visible with a magnify- 
ing-glass). In case. the potassium solution is not very concentrated, 
no precipitation may appear at first; but on rubbing the sides of the 





* Platinic chloride, PtCl, gives no precipitate with potassium salts, or at least 
only after long standing. The above reagent, chloroplatinic acid, is a dibasie acid 
and is obtained by dissolving platinum in aqua regia. The solution is prepared of 
such strength that there are 10 gms. of platinum in every 100 ¢e. 


POTASSIUM | 79 


beaker or test-tube with a glass rod the formation of the precipitate 
will be hastened. 

This is always the case when a crystalline precipitate is formed. 
The solution is supersaturated before the precipitate separates out, 
_ and the formation of crystals is hastened by the mechanical shock. 

The behavior of the potassium chloroplatinate on ignition is charac- 
teristic; it is decomposed into chlorine, platinum, and potassium chloride: 


K2[PtClo] =2KC1+Pt+2Cle 1. 


If the products of ignition are treated with water, and the platinum 
filtered off, the filtrate will again give with chloroplatinic acid the yellow 
erystalline precipitate of K2[PtCle]. (Note difference from ammonium 
chloroplatinate.) 


Solubility of the Potassium Chloroplatinate in Water. 


100 cc. of water dissolve at 0°, 0.70 gm.; 10°, 0.90 gm.; 20°, 1.12 
gms.; and at 100°, 5.18 gms. Ke[PtCle]. 

In a saturated KCl solution, or in 75 per cent alcohol, the pre- 
cipitate is practically insoluble. 

For this reaction it is best to use the chloride. The addition of 
chloroplatinic acid to potassium iodide solution causes a deep-reddish- 
brown color due to the conversion of PtCls~ into non-ionized PtI¢~ - 
ion, of which the potassium salt is more soluble: 


PtCle~ +6I “—> PtI,— +6Cl~. 


Similarly, potassium cyanide is not precipitated by chloroplatinic 
acid, owing to the formation of complex platinum-cyanogen compounds. 

To test an iodide or cyanide for potassium, the salt should first be 
changed to chloride by evaporation with concentrated hydrochloric acid. 

2. Tartaric Acid, H2C4H40¢, produces, in not too dilute neutral 
solutions of potassium salts, a white crystalline precipitate of potas- 
sium acid tartrate (orthorhombic, hemihedral) : 


K*+H2C4Hs40g—-KHCsHi0¢ +H, 


Rubbing the sides of the dish will hasten the formation of the 
precipitate. | 

Potassium acid tartrate is readily soluble in mineral acids, but 
difficultly soluble in acetic acid and water; 100 parts of water at 
10° C. dissolve 0.425 gm. of this salt. If sodium acetate is added to 
the solution, the hydrogen ions set free by the above reaction will 
unite with the acetate ions to form non-ionized acetic acid (cf. p. 46): 


H*+C2H3027 > HC2H302, 





80 REACTIONS OF THE METALS 


whereby the reaction is made much more sensitive. 'Too much sodium 
acetate, however, exerts a solvent action upon the potassium acid tar- 
trate. In the presence of considerable sodium acetate, acetic acid is. 
ionized to such a slight extent that the hydrogen ions from HC4,H40,— 
disappear and, as a result, the potassium acid tartrate dissolves. Neu- 
tralization with caustic alkali is even more dangerous; soluble alee 
tartrate is formed: 


KHC,H.,0,+0OH — Kr+ C4H40, +H20. 


On igniting potassium acid tartrate, empyreumatic vapors (smelling 
like burnt sugar) are given off and a mixture of carbon and potassium 
carbonate is left behind. If the mass is now moistened with hydrochloric 
acid, it will froth strongly. This is a property not only of potassium 
tartrate, but is common to all alkali salts of organic acids. On ignition 
they are changed into carbonates, and when the acid is non-volatile, 
carbonization takes place; but with volatile acids there is at the most 
only a slight carbonization. With many metals the carbonate is not 
left unchanged; frequently it is broken up into carbon dioxide and 
oxide of the metal; in the case of salts of organic acids with reducible 
metals, the metal itself is left with the carbon. Thus sodium acetate 
will yield sodium carbonate and acetone, with only a slight carboniza- 
tion: 
2CH3CO -ONa= NaeCO3 +CHsCOCHs. 


Acetone (colorless 
volatile liquid, 
burns with 
luminous flame) 
On gentle ignition, calcium oxalate yields calcium carbonate and 


carbon monoxide; the latter burns with a blue flame. 


CaC204=CaCO3+COT. 


On strong ignition, the calcium carbonate is decomposed into lime and 
carbon dioxide: 


CaCO; — CaO-+-COs; fT. 


Tartrates of lead, iron, and many other metals on being ignited yield 
carbon and metal. 

3. Bismuth-Sodium Thiosulphate (Carnot’s* reaction).—If one drop 
of half-normal bismuth nitrate solution is mixed with two or three 
drops of half-normal thiosulfate solution and 10-15 cc. of absolute 
alcohol (any turbidity being removed by the careful addition of a very 





*Z. anal. Chem. (1897), 512. 


ras. 











POTASSIUM 81 


little water), a sensitive reagent for potassium ions is prepared in which 
the bismuth is present in a complex anion: 


Bit **+38203" — Bi(S20s)3". 


The sodium salt is soluble in alcohol but the yellow potassium salt is not: 


Bi(S203)3-+3K* — K3[Bi(S203)3). 


The presence of ammonium chloride prevents the reaction. 
4, Fluosilicic Acid, H2Sil’s, added in considerable excess to a solution 
of a potassium salt, precipitates gelatinous potassium fluosilicate, 


Sik’. + QKt = K2SiF¢ ’ 
which is difficultly soluble in water and dilute acids and insoluble in 


alcohol. On heating, it is decomposed into volatile silicon fluoride, 
and potassium fluoride remains behind: 


K.2SiFs=2KF+SiF; 7. 
5. Perchloric Acid, HC1O4, precipitates white, crystalline potassium 


perchlorate, 


HClO. +K+t— Ht+KCl0,; 


100 ce. of water at 0° dissolve 0.07 gm., and at 100° 19.8 gms. KCI1Ou. 
It is so slightly soluble in 97 per cent alcohol that the precipitate can 
be obtained with less than 2 mgs. of potassium ions. 

6. Sodium Cobaltinitrite, Nas3[Co(NOz)«], precipitates yellow 
potassium-sodium cobaltinitrite from neutral or slightly acid solution: 


Nas[Co(NO2)¢] +2KCI=K2Na[Co(NO2e).¢]+2NaCl. 


Ammonium salts give a similar precipitate, but moderate amounts of 


~alkaline-earth elements or of lithium and sodium do not interfere. - 


The test must not be made in alkaline solution or Co(OH)3 will be pre- 
cipitated. The reagent, prepared according to the directions on p. 74, 
permits the detection of 0.3 mg. of potassium within ten minutes. If 
the reagent is prepared according to the following directions of Biil- 
mann,* as little as 0.0009 mg. potassium can be detected in the presence 
of 4000 equivalents of sodium. 


PREPARATION OF SODIUM COBALTINITRITE 


Dissolve 150 gms. of sodium nitrite in 150 cc. of hot water and allow the 
solution to cool to about 40°, which will cause the deposition of some sodium 
nitrite crystals. Add 50 gms. of cobalt nitrate crystals, stir rapidly and add 
50 ce. of 50 per cent acetic acid in small portions; stopper the flask and shake 
vigorously. Pass a rapid stream of air through the liquid and then allow it 


* Z. anal. Chem., 39 (1900), 284. 





82 REACTIONS OF THE METALS 


to stand quietly over night. In the morning, more or less brown precipitate 
will be found on the bottom of the flask due to a little potassium salt present 
in the sodium nitrite. Siphon off the clear liquid through a filter and add to 
the filtrate, while stirring, about 200 cc. of aleohol in small portions; this causes 
the precipitation of the greater part of the dissolved sodium cobaltinitrite. 
After several hours, filter off the precipitate and drain it as completely as possi- 
ble with the aid of suction. Wash the precipitate four times with 25-ce. portions 
of alcohol and twice with ether. Recrystallize the salt by dissolving each 10 
gms. of solid in 15 ec. of water and precipitating with 35 cc. of alcohol. The dry 
salt keeps very well, but the aqueous solution decomposes gradually. To 
obtain the best results, the reagent should be freshly prepared by dissolving 
1 gm. of the salt in 10 ce, of water. . 


REACTIONS IN THE DRY WAY 


Potassium compounds color the non-luminous flame’ violet. The 
presence of very small amounts of sodium obscures the violet color, 
but if the flame is viewed through cobalt glass or indigo solution, the 
reddish-violet potassium rays pass through, while the yellow sodium 
rays are completely absorbed. | 

Flame Spectrum.—Potassium gives a characteristic flame spec- 
trum. A double red line, 769.9 yu and 766.5 yy (appearing as a single 
line with weaker dispersion), and a faint violet line, 404.4 uu, appear 
at comparatively low temperatures. With a hotter flame, other 
lines are visible in the yellow, 583.2 yu; 580.2 wu and 578.2 wy, and in » 
the green, 535.1 wu and 511.3 py (see chart, Frontispiece). 

Gooch and Hart * were able to detect 0.001 mg. of pure potassium 
as chloride in 0.02 cc. water, but the presence of one hundred times 
as much sodium caused the potassium lines to disappear unless the 
sodium rays were deflected from the field. 


SODIUM, Na. At. Wt. 23.00 
Sp. Gr. 0.97. M. Pt. 95.6° C. 


Occurrence.—Sodium occurs very extensively in nature. Its most 
important mineral is halite, rock salt (NaCl), isometric system. 
Halite is found in very large deposits often quite pure, but usually 
contaminated with clay, anhydrite, and gypsum, and is present — 
in large amounts in the ocean, and in many salt springs. Sodium 
also occurs in nature in the form of carbonate, as thermonatrite 
(Na2CO3-H20), orthorhombic; natron or soda (Na2CO3-10H20), 
monoclinic; trona (NazCOz3-NaHCO3-2H2O), monoclinic; as nitrate | 
in Chili saltpetre, or soda nitre (NaNOs3), hexagonal, rhombohedral; 
as eryolite (NagAlF’s), triclinic; in many silicates as albite (NaAISi3Qg), 
triclinic; and as tinkal, borax (NazBsO7-10H20), monoclinic. 


* Z. anal. Chem., 36 (1897), 390. 








SODIUM 83 


REACTIONS IN THE WET WAY 


1. Potassium Pyroantimonate,* KsH2Sbe207, produces in neutral 
or weakly alkaline solutions of sodium salts a heavy, white, crystalline 
precipitate, which is formed more quickly by rubbing the sides of the , 
vessel with a glass rod: 


KeHeSbe07 + 2NaCl =. a2HeSbeO07 + 2KCI. 


The test must not be made in an acid solution, for in that case 
an amorphous precipitate of pyroantimonic acid will be formed: 


KeH2Sb207+2HCl = H4Sb207+2KCl. 


Furthermore, no other metals than the alkalies sheuld be present, 
because they also cause precipitates—amorphous ones for the most 
part. 

2. Tartaric Acid and Chloroplatinic Acid do not precipitate 
sodium salts, the sodium ‘salts of these acids being soluble in alco- 
hol as well as in water. (Note difference from potassium.) Sodium 
chloroplatinate is orange in color. 

3. Hydrochloric Acid and Alcohol precipitate sodium chloride. 
Sodium chloride is prepared pure for chemical purposes by passing 
hydrogen chloride gas into the saturated aqueous solution of the salt 
and expelling the moisture and hydrochloric acid from the crystals 
by heating them. If dry hydrogen chloride is passed into an alcoholic 
solution of a sodium salt, less than 1 mg. of sodium will remain in 
solution. | 


Sodium Peroxide, NazO2 


This substance, which is now used commercially on account of 
its energetic oxidizing power, is obtained as a heavy, yellow powder, 
by burning dry sodium in the air; it shows the following character- 
istic reactions: 

Behavior toward Water.—If a little water is added to some of this 
substance in a test-tube, considerable heat is evolved and oxygen gas 
is liberated (sufficient to ignite a glowing splinter ¢). Water decom- 
poses the sodium peroxide, according to the equation 


Naz02+2H20 =2Na0H+Ho20o. 





* For the preparation of this reagent see page 74 and under Antimony. 

+ This will sometimes cause an explosion. Commercial sodium peroxide often 
contains metallic sodium, which with water forms hydrogen; thus both hydrogen 
and oxygen are set free at the same time, and the glowing splinter may then cause 
an explosion. (Private communication from E. Constam.) 


84 REACTIONS OF THE METALS 


But on account of the heat of the reaction a part of the hydrogen — 
peroxide is decomposed into water and oxygen. 

If the solution is kept cold, which can be done by throwing the 
sodium peroxide in small portions into ice-water, it will dissolve with 
scarcely any evolution of oxygen, to a clear, strongly alkaline liquid, 
which gives, as before, all the reactions of hydrogen peroxide. 

If some sodium peroxide is placed on a watch-glass under a bell- 
jar and near an evaporating-dish containing water, the sodium per- 
oxide in twelve hours will completely change over to a pure white 
hydrate (NagO2+8H20), which will dissolve in water without decom- 
position at the ordinary temperature. By standing in a desiccator 
over sulfuric acid, the octohydrate is changed to NazO2+2H:20. 

> 


Reactions of Hydrogen Peroxide 


(a) In Acid Solution 


If the solution obtained by the action of water on sodium per- 
oxide is used for these tests, it must be acidified with dilute sulfuric 
acid, care being taken to keaists the solution cool. 

1. Titanium Sulfate gives a distinct yellow color, caused by 
the formation of pertitanic acid, 


4 
hy t4 HeOo +. 2HeO <> 4H HeTiOs 


This is the most delicate test for hydrogen peroxide. The tita- 
nium sulfate solution for this reaction may be prepared by fusing 
one part of commercial titanium dioxide with 15-20 parts of potassium 
pyrosulfate and dissolving the fusion, after cooling, in cold, dilute 
sulfuric acid. It may also be prepared by heating titanium dioxide 
with concentrated sulfuric acid, until a clear solution is obtained, 
cooling and diluting carefully i water. 

The addition of caustic alkali, ammonia or ammonium carbonate 
gives a yellowish-orange precipitate which redissolves in an excess of the 
reagent. Classen has used this reaction as a method for separating 
titanium from ferric iron.* Some chemists prefer to write the formula 
of pertitanic acid as TiO2-H2O2, which assumes a true peroxide 
structure instead of hexivalent titanium. Possibly a condition of 
equilibrium exists between the two structures: 


ou ee O-.OH 
ae > O= TS 


* Ber., 21 (1888), 370. 








“3 
J 

; 

i 

> 
a 
J 
: 
y. 
A 


” 


SODIUM | 85 


2. Chromic Acid.—If the acid solution of hydrogen peroxide is 
shaken with a little ether (free from alcohol) and a trace of potassium 
dichromate is added, after which the mixture is again shaken, the 
upper layer of ethereal solution will be colored a beautiful blue, owing 
to the formation of chromium peroxide (cf. p. 18). 

This test is very sensitive and will detect as little as one-tenth 
milligram of H2O2. In carrying out this test, a blank test must always 
be tried with the ether and alcohol alone, because the former will often 
give the test. Ether, after standing in the air, is likely to contain 
some ethyl peroxide (C2H5)403, (?) which gives the test. It is possible 
to free the ether from this peroxide by letting it stand over night 
in contact with sodium and then redistilling it. 

3. Permanganic Acid in acid solution will be decolorized, with 
evolution of oxygen: 


2Mn04-+5H202+6Ht > 2Mn+++8H20+50> 1. 


Similar to the permanganate, many other oxides are reduced by 


hydrogen peroxide, with evolution of oxygen; e.g., AgeO, Pb3Qa, 
PbO2, MnOzg, Co20s, ete.: 


Ago0+H202— H20+02 T +2Ag 
Mn02+H202+2H* > Mn**+2H20+02 T 
Co203+H202+4H* — 2Cott+3H20+02 1. 


4. Potassium Ferricyanide and Ferric Chloride.—If a trace of 
potassium ferricyanide is added to a very dilute and nearly neutral 
solution of ferric chloride, so that the solution appears a distinct 
yellow, and a nearly neutral solution of hydrogen peroxide is then 
added, the mixture will soon assume a green tint, and finally, on stand- 
ing, Prussian blue will separate out. The potassium ferricyanide 
is reduced by the hydrogen peroxide to potassium ferrocyanide, 
which forms Prussian blue with the ferric chloride. 


2Fe(CN)6~ +H202 — 2Fe(CN)g--+02 T +2H* 
and 3Fe(CN)g _+4Fet** —  Fea[Fe(CN)¢]3. 


According to Schénbein, as little as 0.02 mg. H2Oz per liter may be 
detected by this reaction. As many other substances (SnCle, SOs, 


etc.) will reduce potassium ferricyanide to potassium ferrocyanide, 


this reaction alone is not always a reliable test. 

5. Starch Paste and Potassium Iodide.—If to an acid solution 
containing starch paste and potassium iodide some hydrogen peroxide 
is added, a blue color will at once appear: 


2K1+H202=2KOH-+Iz. 


86 REACTIONS OF THE METALS 


By means of this reaction, 0.05 mg. per liter of hydrogen peroxide may 
be detected. 


(b) In Alkaline Solution 


1. Gold Chloride by means of hydrogen peroxide at ordinary 
temperatures will be reduced to metal, with evolution of oxygen. 
The gold usually separates in a very finely divided state, and appears 
brown by reflected light and greenish blue by transmitted light: 


2Autt+++3H.0.+60H~ > 2Au+6H20+302 fT. 


If very dilute gold solutions are used, the metal sometimes separates 
out in the form of a yellowish film adhering to the sides of the test-tube. 
2. Salts of Manganese and Cobalt give dark-colored precipitates: 
Mntt+20H~+H:202 — H20+Mn0O(0H)>2; 


Brown 


2Cot *+20H +H202 — H20+2Co(OH)3. 
Black 
Hypochlorites give the same reactions with manganese and cobalt 
salts, but they do not give the reaction with gold chloride. 


Ozone, O3 


Ozone is always formed when oxygen is exposed to the silent electric 
discharge. It is often present in oxygen that has been prepared elec- 
trolytically and, according to Brunck, is present to:some extent in the 
gas prepared by ignition of potassium chlorate. Ozone is a strong 
oxidizing agent and behaves in many respects like hydrogen peroxide, 
with which it is often confused. Ozone may be distinguished from 
hydrogen peroxide as follows: 

1. Ozone does not give a yellow coloration with titanium sulfate 
solution. . 

2. Ozone does not cause precipitation of gold from its solutions. 

3. Ozone sets free iodine immediately from dilute, neutral potassium 
iodide solution. 

4. Ozone liberates bromine from an acid solution of sodium bromide. 

5. Ozone causes bright metallic silver at once to assume a steel- 
blue tint. 

The sensitiveness of this last reaction is remarkable if carried out 
according to the directions of Manchot and Kampeschulte. Heat a 
bright piece of silver foil to about 240° and then expose it to the action 
of ozone; steel-blue spots with violet edges at once appear. This 
reaction does not take place with pure silver in the cold. If, however, 
the silver is polished by rubbing with emery paper, the reaction will 





r Pinta al > 
vy at 


AMMONIUM | 87 


, then take place in the cold. Traces of iron oxide are left upon the 
silver from the emery and catalyze the reaction. Other oxides, AgoO, 


CoO, NiO, BizO3, PbsO04, -V205s, MnOz, CuO, ThO2, CeOe, TiOse, 
WO3, U30s, and to a less degree MoOz, HgO, CaO and BaO, have a 
similar effect. Thus if the silver is etched with nitric acid and is then 


_ dried, it will react with ozone in the cold. 


The principal reactions of sodium are the 


REACTIONS IN THE DRY WAY 


Sodium salts color the non-luminous gas-flame a monochromatic 


_ yellow, which can be readily distinguished from the yellow flame of 
the gas in the following way: If we illuminate an orange-colored body 


(such as a stick of sealing-wax or a crystal of potassium dichromate) 
with white light (all glowing solid bodies emit white light), the red and 
orange rays will be reflected: the body appears orange. If these bodies 
are illuminated with the monochromatic sodium light, they can now 
only reflect yellow light: the bodies appear yellow (a delicate test). 

Flame Spectrum.—A yellow double line (589.6 wu and 587.0), 
coinciding with the D-line of the sun’s spectrum. This is an extremely 
delicate reaction; 110-7 mg. of sodium can be recognized in the spec- 
trum. 

AMMONIUM, NHy. At. Wt. 18.04 


Occurrence.—In small amounts as carbonate and nitrite in the 
air; as ammonium chloride in the fissures of active volcanoes. Am- 


‘monium derivatives are formed by the decay of many organic substances 


containing nitrogen: albumin, urea, etc., 
CO(NHz)2+H20 =C02+2NHs, 


and in a similar way by the dry distillation of many nitrogenous sub- 
stances, such as coal, horn, hair, etc. 

Although ammonium itself is known only in the form of its amalgam, 
we are justified in considering it as a metal, in the first place because the 
electrolysis of ammonium salts causes the setting free of the cation 
NH,4(NH3+H) at the same time. that the corresponding anion is set 
free; and, further, because the ammonium salts are isomorphous with 
potassium salts. 


REACTIONS IN THE WET WAY 


ra: Strong bases, NaOH, KOH or Ca(OH)s, added to an ammonium 
salt in the presence of a little water cause the evolution of ammonia 
on heating; the gas can be recognized by its odor, by fumes of 
ammonium chloride being formed when a rod moistened with 12 N 


88 REACTIONS OF THE METALS 


hydrochloric acid is placed in contact with the vapors, by its turning 
red litmus blue; or by the blackening of mercurous nitrate paper: 


Hg 
2Hg2(NOs)2-+4NHs-+H20=3NHiNOs-+0C NH2-NO3+2Hg 
Hg 


/ 





Black 


The reaction of strong bases upon ammonium salts may be explained on 
the basis of the laws of chemical equilibrium. Ammonia, NH,, is a gas which 
is very soluble in cold water and insoluble in boiling water. One volume of 
water at 0° and 760 mm. dissolves 1800 volumes of the gas; at 20° it dissolves 
710 volumes; at 100° all the gas can be expelled, there being no constant boil- 
ing mixture as in the case of hydrochloric acid (p. 57). 

The solution of ammonia in water at the laboratory temperature is in 
equilibrium with ammonium hydroxide, 


NH;+H:0 — NH,OH, 


and, for this equilibrium, the mass action expression (p. 14) is 
(N H;|X [HO] 
[NH.OH] 


In a dilute solution of ammonia, the absolute quantity of water present is not 
changed much as a result of this equilibrium, so that when an accuracy corre- 
sponding to only two significant figures is desired, the H,O member of this 
expression may be regarded as a constant (cf. p. 48) and the equilibrium 
expression then becomes 


=k. 





[NH3] _ . 
(NH,OH] =kyus (I) 
The value of this constant has been found to be about 2 at 20°. The 


ammonium hydroxide, however, is not only in equilibrium with ammonia, 
but also with ammonium and hydroxy] ions, 


NH,OH — NH,t+0H-, 
and the mass-action expression of this equilibrium is 
[NH,*] [OH] 

[NH,OH] 


Moore * found the value of this constant to be about 5X10—. By adding 
[NH,0H] 





= kpase (II) ° 


to the left-hand member of equation (I) and its equivalent 1 to the 








[NH,OH] 
right-hand member, we obtain the equation 
[NH;]+[NH,OH] _ pee 
vH;|+[NH,OH 
and [NH,oH) =#! a tit) 





* J. Chem. Soc., 91 (1907), 1379. 





AMMONIUM 89 


Inserting this value for [NH,OH] in (II) and transferring k’ to the right-hand 
member, we have 

| [NHAOH=]__ frase _ 7 
[NH.OH]+[NH;] k 9” 


The value of this constant, K, at 18° is about 1.8107. 

The ammonium salts, unlike the free base, are largely ionized. When, 
therefore, an excess of OH is added to the solution of an ammonium salt, 
it is necessary, in order to establish equilibrium between NH,* and OH™ for 
the greater part of the NH,* to be converted into NH,OH and then, to establish 
equilibrium between NH,OH and NHs, about two-thirds of the NH,OH is 

changed into NH;. By boiling the solution, the NH; is expelled and the 
above-mentioned states of equilibrium-are disturbed and, as the final result, 
all of the original NH,t becomes converted into NH; gas. Less than 0.2 mg. 
of ammonium can be detected by the litmus test when properly carried out. 
Care should be taken not to boil the solution so that some of the alkaline liquid 
becomes spattered into the nostril or upon the test paper. 








Certain complex ammonia derivatives do not always evolve ammonia 


aes 2 ; /NEib . 
in this test. When pure mercuric amidochloride, Hg , 1s heated 


Nc 

with caustic soda solution, a part of the nitrogen is evolved as ammonia, 
but if considerable mercuric salt is present the test is not obtained. 
This is because the mercuric amidochloride itself is only slightly soluble 
in water and, especially in the presence of an excess of mercuric. com- 
pounds, furnishes scarcely any ammonium ions. If some potassium 
sulfide is added, however, the mercury is converted into more insoluble 
mercuric sulfide and the ammonia test is obtained: 


NH, 
HeC +K>8+H»0 =HgS+KCl+KOH+NH; f. 
Cl 


Water itself in some cases causes evolution of ammonia gas. It 
decomposes many nitrides, metal amides and cyanamides: 


Mg3N2+6HOH =38Mg(OH)2+2NH3 17, 
NH2Na+HOH =NaOH+NHs317,. 
CaCN2+3H20 = CaCO3+2NHs3 ¢ (at high temperatures) 


Calcium 
cyanamide 


2. Chloroplatinic Acid gives a yellow crystalline precipitate: 
H2[PtCle]++2NH4* > (NH4)2PtCle+2H". 


This salt may be distinguished from the potassium salt— 
(a) by its behavior on ignition; platinum alone is left behind: 


(NH4)2[PtCle] =2NH4Cl+2Cle T +Pt; 


90 REACTIONS OF THE METALS 


(b) by its behavior on treating with strong bases, whereby the smell 
of ammonia may be detected: 


(NH4)o[PtCle]-+2Na0H = Nas[PtCle]+2H20+2NHz3 Tf. 


3. Tartaric Acid produces, as with potassium, a white, crystalline ~ 
precipitate of ammonium acid tartrate. The addition of a little sodium 
acetate, and rubbing the sides of the glass vessel with a stirring-rod, will 
hasten the formation of the precipitate: 3 


HoC4H406.+NH;t > NH.HC.Hi0,¢+H". 


The ammonium acid tartrate, like the corresponding potassium 
salt, is soluble in alkalies and mineral acids. It may be distinguished 
from the potassium salt by its behavior on ignition, carbon alone 
being left behind, and the residue not effervescing with hydrochloric 
acid; furthermore, ammonium acid tartrate will give off ammonia on - 
Cine heated with caustic soda solution. 

4. Sodium Cobaltinitrite gives a yellow pr ecipitate similar to 
that produced with potassium. Before testing for potassium with 
this reagent, therefore, it is necessary to expel ammonium salts by 
evaporating the solution to dryness in a porcelain dish and heating 
until no more fumes are evolved. 


The above-mentioned reactions are not suitable for the detection of the 
very small amounts of ammonia or of ammonium ions that are found in drinking- 
water. Insuch cases Nessler’s reagent is used (an alkaline solution of potassium 
mercuric iodide). Large amounts of ammonia produce a brown precipitate, 


The precipitate is the iodide of the so-called Millon’s base; its structural 
Hg 
formula is probably 0< Nt -I. The corresponding black nitrate is formed 
H 


in the test for ammonia with mercurous nitrate paper. . The iodide of Millon’s 
base has such a remarkable coloring power that mere traces of the ammonium 
ion can be detected by the yellow or brown color imparted to the solution. The 
test is obtained with ordinary distilled water. Since ammonia is often present 
in water as a result of its contact with decaying organic matter, the test for 
ammonia helps to determine whether a water is suitable for drinking. The test 
is called the Nessler test, and the alkaline solution of potassium mercuric iodide 
is called Nessler’s reagent. 

Water free from ammonia should be used in carrying out the Nessler test. 
Some sodium carbonate and a little potassium permanganate is added to ordi- 
nary distilled water, which is then redistilled, rejecting the first fourth and last 
sixth of the distillate; the middle portion is the so-called best water of the 
chemical laboratory. For the most accurate work, Nessler’s reagent should 


AMMONIUM 91 


be prepared with such water (cf. p. 73) and the test should be made in a labora- 


tory from which ammonium fumes are absent. 





To test a water for the presence of traces of ammonium ions, the apparatus 
shown in Fig. 4 may be used. First of all, the apparatus itself must be freed 
from all traces of ammonium salt. To accomplish this, place about 50 cc. of 
water in the retort, add 1 ce. of a boiled, saturated solution of sodium carbonate 
and distill with the neck of the retort introduced well into the condenser tube. 
It is advisable not to use a rubber connection between the retort and the con- 
denser; the condensed water serves to make a sufficiently tight connection. 

Continue distilling until 50 c¢. of the distillate placed in a white glass 
graduate, or in a so-called Nessler tube, and treated with 1 cc. of Nessler reagent, 





YA 
) Sarace 








A) 














| | | | 
H 
EI 








Fig. 4. 


will show no sign of color after standing five minutes. The apparatus is then 
ready for the test. 

Empty the retort and refill it with 500 cc. of the water to be tested, add 1 cc. 
of the saturated sodium carbonate solution, distill and collect the first 50 cc. 
of distillate. If as much as 0.2 mg. of ammonium is present, a distinct precipi- 
tate forms on adding 1 cc. of Nessler’s solution and stirring; a pronounced 
yellow color is obtained with much smaller quantities. If mere traces of 
ammonium are present, the yellow color appears only on standing. By com- 
paring the depth of color with that similarly produced with known quantities 
of ammonium chloride, a very close estimate of the exact quantity of ammo- 
nium present can be made, 


REACTIONS IN THE DRY WAY 


All ammonium salts are relatively unstable compounds, the degree 
of stability depending, in general, upon the strength of the acid which 
is combined with the ammonium. The carbonate decomposes appre- 
ciably at ordinary temperatures and when exposed to the air gradually 


92 REACTIONS OF THE METALS 


. disappears as ammonia, carbon dioxide and water. Heating in a closed 
tube causes the decomposition of all ammonium salts and either 
ammonia or some other volatile nitrogen compound escapes. 

If the acid is volatile at the decomposition temperature, the whole 
salt is volatilized, often without first melting, and when the vapors 
of ammonia and acid are cooled the solid again forms. This explains 
the ammonium fumes that result when ammonium salts are expelled 
from a solid residue obtained by evaporation of a solution, and it 
explains the sublimate formed when ammonium chloride is heated in 
a closed tube. The acids which form salts that are not volatilized are 
boric, phosphoric, chromic, molybdic, tungstic and vanadic acids. 

It must be remembered, however, that ammonia with its negative 
valence of three contains nitrogen in its lowest state of oxidation. When 
the decomposition of the original ammonium salt takes place, therefore, 
there is often an oxidation of the nitrogen. Thus the decomposition of 
the nitrate results in the formation of nitrous oxide, N2O, and the 
decomposition of the nitrite, sulfate and dichromate yields nitrogen gas. 

The closed tube reactions of typical ammonium compounds may be 
expressed by the following equations: 


NH,Cl = NH3+HCl, 
NHiNO;=2H20+N;0, 
NHiNO2=2H20-+Np, 
3(NH1)2804 = No-+4NH3+6H20+3802, 
(NH) 2C204=2NH3+H20+CO+COr. 


Toward the last some dicyanogen, (CN)s, is formed from the oxalate. 


NaNH4HPO, : 4H2O => NH3 +5H»eO +NaPOs, 
2(NH1)3PO,=6NH3+2H20+2HPO3, 
(NH4)2Cr207 =4H20+Ne2+CreOsz. 


In this last reaction, the chromic oxide remains as a voluminous mass 
looking something like green tea. A realistic voleano effect can be 
obtained by making a mound of ammonium dichromate around a piece 
of paper and then setting fire to the latter. 

Ammonium salts do not impart a characteristic color to the flame; 
the border of the flame is tinged slightly greenish. 





MAGNESIUM 93 


MAGNESIUM, Mg. At. Wt. 24.32 
Sp. Gr. =1.75. M. Pt. =632.6° C. 


Occurrence—Magnesium compounds are found very abundantly 
in nature. The most important minerals are magnesite, MgCOs, 
rhombohedral, isomorphous with calcite; dolomite, (Ca,Mg)CO3; 
brucite, Mg(OH)2, rhombohedral; carnallite, KMgCls+6H20, ortho- 
rhombic; kieserite, MgSO4+HeO, monoclinic; epsomite, MgSO, 
+7H20, orthorhombic; spinel, MgAloOx4, isometric, isomorphous with 
magnetite, Fe3O04, and with chromite, FeCreO4. Magnesium also occurs 
in a great many silicates. Thus almost all the minerals of the olivine 
group contain more or less magnesium. To this group belong forsterite, 
Mg2S8i04; monticellite, CaMgSi0O4; and olivine, FeMgSiO«z; all ortho- 
rhombic. An important decomposition product of the olivine minerals 
is serpentine, Mg3H4Si2O9. Almost all the minerals of the pyroxene- 
amphibole group, which are all related to orthorhombic enstatite, 
MgSi03, contain magnesium: augite, MgAloSiOg; hornblende, an 
isomorphous mixture of Mg3CaSi4Oi2 and 2(MgAl2Si0¢) ; and tremolite, 
CaMg3SisO12, all three forming monoclinic crystals. Asbestos is a 
variety of tremolite with very fine fibers. Meerschaum is a magnesium 
silicate of the composition HsMgeSisOi0, and is quite similar to tale, 
H2Mg3SisO12, sometimes called steatite. Magnesium also occurs in 
the vegetable kingdom, being an essential constituent of the complex 
organic compound chlorophyll. 

Properties of Magnesium.—Magnesium is a silver-white metal. 
It decomposes water very slowly, forming an oxide, MgO, which 
is only slightly soluble in water, forming magnesium hydroxide; the 


small quantity that dissolves is largely ionized, so that the solution has 


a faint alkaline reaction. Magnesium reacts directly with nitrogen at 
300° C., forming magnesium nitride (Mg3N2), which is readily decom- 
posed by water, forming magnesium hydroxide and ammonia: 


Mg3N2-++6HOH =3Mg(OH)2+2NHs3 fT . 


The salts of magnesium are almost all colorless and soluble in water. 
The solubility products of magnesium hydroxide, carbonate, phosphate, 
arsenate and arsenite are so small that these substances may be regarded 
as insoluble. The sulfide, which can be prepared only in the dry way, 
is completely decomposed by water into hydroxide and hydrogen sul- 
fide (hydrolysis). If an aqueous solution of magnesium chloride be 
evaporated to dryness on the water-bath, there is no hydrolytic decom- 
position, the residual salt, MgCl2+6H20O, dissolves in water, forming a 
perfectly clear solution. On heating the chloride containing the water 


94 REACTIONS OF THE METALS 
fx 


of crystallization to 106° and higher, however, a considerable amount 
of hydrochloric acid fumes escape and a basic magnesium chloride 
insoluble in water is left behind: 


Cl 
MENS 
2MgCle +He2O —_ O +2HCIl. 
Me 
Nol 


When a saturated solution of magnesium chloride is mixed with 
magnesium oxide, the mixture soon solidifies, forming a mass hard 
as stone, known as magnesia cement, consisting of basic magnesium 
chloride. 


REACTIONS IN THE WET WAY 


1. Strong Bases, such as the soluble hydroxides of sodium, potassium 
and barium, precipitate white, gelatinous magnesium hydroxide, and the 
precipitation is practically complete in the absence of ammonium salts 
or if the ammonium salt is all decomposed by boiling with an excess of 
the strong base (cf. pp. 19, 46). 


Mgt ++20H- — Mg(OH)2 


The solubility product of magnesium hydroxide (cf. p. 22) is about 3.4107" 
at the laboratory temperature. The saturated solution of magnesium hydrox- 
ide in pure water contains about 0.0002 mole or 0.012 gm. Mg(OH), per liter. 
In the presence of an excess of OH, the solubility of the magnesium hydroxide 
is much less, as a result of the common ion effect (cf. p. 45) and it is possible, 
by keeping the volume of the solution small and using a slight excess of the 
reagent, to leave less than 1 mg. of magnesium in solution. 

The precipitation of magnesium hydroxide by means of the slightly ionized 
ammonium hydroxide can never be made complete, and if the solution already 
contains ammonium ions in sufficient excess, no precipitation of magnesium 
hydroxide takes place. Moreover, if a precipitate of magnesium hydroxide is 
boiled with a solution of an ammonium salt, such as ammonium chloride, the 
precipitate dissolves. This behavior is due to the fact that the ionization of 
ammonium hydroxide (cf. p. 19 and p. 88) is repressed to such an extent, 
as a result of the common ion effect, that not enough OH~ ions are present at 
any one time to satisfy the solubility product of magnesium hydroxide, and 
even the OH~ from the Mg(OH),. must be in equilibrium with the NH,* ions 
from NH,Cl. 

Formerly, the non-precipitation of magnesium by ammonium hydroxide 
was explained by the assumption that complex salts such as NH.{[MgCl,] or 
(NH,).[MgCl,] were formed, but this explanation has proved untenable.* 





* Cf. Loven, Z. anorg. Chem., 11 (1896), 404; TreapweE.., ibid., 37 (1903), 326 
and Herz, ibid., 38 (1903), 138. 





MAGNESIUM 95 


_ 2. Mercuric Oxide heated with solid magnesium chloride converts 
the latter into magnesium oxide which does not dissolve appreciably in 
water. The mercuric chloride formed and the excess of mercuric oxide 
are volatilized. 

3. Ammonium Carbonate precipitates, in the absence of other 
ammonium salts, a basic salt (usually only on boiling or after long 
standing). The composition of the precipitated salt varies with the 
temperature and the concentration of the solution, the following salt 
being often obtained: 


4Megtt +4(NH4)2CO3 +H20 — Mg, (CO3)3(OH)2+CO2T +8NH4,", 


The addition of an excess of ammonium carbonate reagent (p. 72) 
and an equal volume of 95 per cent alcohol causes the complete pre- 
cipitation of magnesium as MgCO3- (NH4)2CO3-4H20 from a cold, con- 
centrated solution of magnesium salt. 

The magnesium ammonium carbonate is fairly soluble in water 
and the solubility increases rapidly with rise of temperature. Thus 
no precipitate is obtained upon the addition of ammonium carbonate 
to a hot dilute solution of magnesium salt containing ammonium 
chloride and no alcohol. (Note difference from barium, strontium 
and calcium.) | 

' 4. Sodium Phosphate is the characteristic reagent for magnesium. 
It produces in solutions containing ammonium chloride, and in the 
presence of ammonia, a white crystalline precipitate (orthorhombic, 
hemimorphous) of magnesium-ammonium phosphate, * 


Mgt Fo NH4t+ POw— MgNH,PO,. 


From very dilute solutions the precipitate separates only after 
standing some time, owing to the tendency to form supersaturated solu- 
tions. Rubbing the sides of the beaker with a glass rod hastens the 
formation of the precipitate. 


The table on p. 22 gives the solubility product of magnesium ammonium 
phosphate as 2.5X107~; about 0.0086 mg. of the salt dissolves in a liter of 
water at the laboratory temperature. The solubility in’water is increased by 


its tendency to undergo hydrolysis. 
MgNH,P0.,+HOH @ Mgt *++HPO~+NH.OH. (I) 


The tendency to undergo hydrolysis increases rapidly with rise in temperature 
(cf. p. 51). A similar decomposition is caused by hydrogen ions alone, 


MegNH,PO,+Ht @ Met++HPOC+NH.t, (II) 





* Magnesium ammonium phosphate crystallizes with six molecules of water. 


96 REACTIONS OF THE METALS 


which lessen the tendency for HPO, to ionize; the precipitate, therefore, dis- 
solves readily in the presence of any acid which is ionized more than HPO,-, even 
acetic acid (cf. p. 10). The presence of ammonium hydroxide prevents the 
hydrolysis, in accordance with the mass-action law. For this reason an excess 
of ammonia solution is usually added. Ammonium chloride, by virtue of the 
common ion effect, lessens the quantity of Mgt* ions required to reach the 
solubility product and causes reaction (II) to take place in the direction right 
to left; but, on the other hand, it should favor reaction (I) somewhat, because 
it represses the ionization of ammonium hydroxide. As a matter of fact 
ammonium salts usually retard the formation of the precipitate, but do not 
eventually make it more soluble if an excess of ammonia is present. 

Neubauer * has shown that the conditions are still more complicated on ac- 
count of the tendency for gelatinous Mg;(PO,). and crystalline Mg(NH,)4(PO,)2 
toform. Tertiary magnesium phosphate, Mg;(PO,)s, is formed in cold, strongly 
ammoniacal solutions containing but little ammonium salts. The mono- 
magnesium-tetrammonium phosphate is formed in neutral or slightly alkaline 
solutions containing considerable ammonium salts. B. Schmitz? has shown 
that beautifully crystalline precipitates can be obtained in the presence of 
ammonium salts by adding sodium or ammonium phosphate to the boiling, 
acid solution of the magnesium salt. Then, on adding one-third the solution’s 
volume of 6 N ammonia and allowing the solution to cool, complete precipi- 
tation, as MgNH,PO,-6H,0, takes place. 


REACTIONS IN THE DRY WAY 


All magnesium salts are more or less changed on heating in the 
air, leaving behind the oxide or an insoluble basic salt. On char- 
coal with sodium carbonate before the blowpipe, magnesium com- 
pounds are changed to white magnesium oxide, which is strongly 
luminous when hot. Calcium, strontium, and aluminium compounds 
behave the same way. The magnesium salts are nonvolatile, do not 
color the flame, and give no flame spectrum, but do give a characteristic 
spark spectrum. 


Detection and Separation of Magnesium and the Alkalies in the Pres- 
ence of One Another 


All of the typical reactions that have been described for magnesium 
can be used for separating magnesium from the alkalies. The test 
for ammonium is always carried out with some of the original substance 
or solution. | 

Ordinarily in qualitative analysis, it is customary either to test for 
magnesium in one portion of the solution by means of sodium phosphate 
and for sodium and potassium in another portion, or to precipitate the 





* Z. angew. Chem. (1896), 439. Cf. Goocn and Austin, Z. anorg. Chem., 20, 121. 
+ Z. anal. Chem. (1906), 512. . Cf. JonaEnsEN, ibid. (1906), 278. 


MAGNESIUM 97 
magnesium as carbonate and test for sodium and potassium in the 
filtrate. In this last method the magnesium can be precipitated with 
the alkaline earths. If this is done, the test for magnesium is not made 
in the analysis of this group. 

In all the principal systematic methods given in this book for the 
examination of the basic constituents, a very brief tabular outline 
will be given first with reference by numbers to the detailed direc- 
tions that immediately follow. 


TaBLE I.—ANALYSIS OF Group V. Mertuop A. 




















NH;+. Test a|Mgt*t. Test a| Evaporate to dryness. Expel NH, salts. 
portion of the} part ofthe fil- Dissolve in water, add Ba(OH)2 and reject 
original sub-| trate from| _ precipitate. Add. HCl, NH,OH and 
stance for| Groupll/V, (NH4)2CO; and also reject this precipitate. 
NH; by boil-| Method A, for Evaporate, expel NH, salts, dissolve in water, 
ting with| Mg with| filter, and add HCl0O,4. Evaporate, add 
NaOH. (1) NH,OH_ and alcohol and filter. (8). 

N. a2H P O14. (2) 

Precipitate: KClO,.| Filtrate: Na, Satu- 
Dissolve in hot| rate with HCl gas. 
water and add| Filter off NaCl and 
NasCo(NOz2)eé.| reject filtrate. Dis- 
Yellow precipitate is | solve in water and add 
K.,NaCo(NOz2)6. (4)| KeH2Sb.07. Crystal- 
line precipitate is 

NazH.Sb2O7. (5) 

PROCEDURE 


1. Test for Ammonium. Place a little of the original substance (correspond- 
ing to about 0.25 gm. of solid) in a test-tube, add about 2 cc. of 6 N sodium 
hydroxide, heat nearly to boiling and hold a piece of red litmus paper, wrapped 
around the end of a stirring rod, in the escaping vapors. Take care not to allow 
any of the caustic alkali to come in contact with the litmus either by spattering 
or by allowing the paper to touch the sides of the test-tube. A good idea of the 
quantity of ammonium present can be estimated by the odor. If it is desired 
to test for traces of ammonium, carry out the Nessler test as described on p. 91. 

2. Test for Magnesiuin. Dissolve the substance in as little water as possible, 
or, if a solution, evaporate to dryness, moisten with 6 N hydrochloric acid, 
heat gently, dissolve in a little water and filter if necessary. The addition of the 
acid is not usually necessary, but sometimes, when calcium has previously been 
removed as oxalate, a difficultly soluble oxalate of magnesium and ammonium 
is formed which is best dissolved by treatment with acid before adding water. 
If the solution to be tested contains considerable ammonium salt, it is usually 
best to expel it by igniting the residue obtained by evaporation; the treatment 
of the residue with hydrochloric acid is then absolutely necessary, in order to 
decompose basic magnesium salts which are formed during the ignition. 

If ammonium chloride is not already present, add half the solution’s volume 
of normal ammonium chloride and enough ammonia to make the solution dis- 


98 REACTIONS OF THE METALS 


tinctly ammoniacal. If a precipitate of magnesium hydroxide is formed, add 
more ammonium chloride to dissolve it. A flocculent precipitate produced at 
this point may be aluminium hydroxide or silica. Such a precipitate should 
be filtered off and discarded. To the clear ammoniacal solution add a little 
sodium phosphate solution and rub, with a rounded stirring rod, the inside 
walls of the vessel containing the solution. If as much as a few tenths of a 
milligram of magnesium is present per 100 cc., a crystalline precipitate of 
magnesium ammonium phosphate should form within a few minutes. In case 
no precipitate is noticeable, set the beaker aside and allow it to stand overnight. 
Traces of magnesium phosphate will usually form on the sides of the beaker 
where it was scratched by the stirring rod. 

If the precipitate does not appear distinctly crystalline, it may contain 
aluminium phosphate. In such cases filter off the precipitate and redissolve 
it in 6 N acetic acid, which will not dissolve aluminium phosphate, but 
readily dissolves magnesium ammonium phosphate. Heat the solution nearly 
to boiling, add a liberal excess of ammonia and allow to cool; a crystalline pre- 
cipitate should be obtained by this treatment. 

3. Test for Sodium and Potassium. Before testing for sodium and potassium 
it is best to remove the magnesium. Evaporate the solution to dryness in a 
porcelain or platinum dish and gently ignite the residue to expel ammonium 
salts. Dissolve the residue in a little water and, without paying any atten- 
tion to any residue of basic magnesium salt, add barium hydroxide solu- 
tion until strongly alkaline. Heat to boiling and filter off the precipitate, which 
may contain magnesium hydroxide, barium sulfate, if any sulfate ions were 
present, and barium carbonate from the contact of the barium hydroxide 
solution with the air. These operations should never be carried out in glass 
dishes because of the danger of obtaining alkali from the glass. Make the fil- 
trate from the magnesium precipitate barely acid with hydrochloric acid and 
remove all barium ions by the addition of ammonia and ammonium earbon- 
ate. Filter off the precipitated barium carbonate, evaporate the filtrate to 
dryness and expel the ammonium salts by careful ignition. Dissolve the resi- 
due in a little water and make sure that all the barium was removed by add- 
ing a little more ammonium carbonate. Filter if necessary and again evapo- 
rate to dryness. Expel the ammonium salt by ignition, cool and wash down 
the sides of the dish with about 5 cc. of water to dissolve any ammonium salt 
that may possibly be left there. Evaporate to dryness again and expel the last 
traces of ammonium salts by igniting at a dull red heat until no more fumes 
are evolved. Dissolve the residue in a little water and filter the resulting 
solution through a small filter. A black residue is often due to carbonization 
of pyridine bases, which are commonly present to a slight extent in ammonia 
solutions. 

Evaporate the solution just to dryness, add 5 to 15 ce. of 2 N perchloric 
acid and evaporate carefully, by keeping the dish in constant motion over a 
free flame, until dense fumes of perchloric acid are evolved. Cool completely 
and add 20 ce. of alcohol. . 

The perchloric acid solution must not be heated or concentrated by evaporation 
after the alcohol has been added or a dangerous explosion is likely to result. 

Stir the solution and press down gently with a stirring rod upon any crystals 
that may be present. If after stirring a few minutes a residue of potassium 
perchlorate remains, add 3 cc.-more of the perchloric acid reagent. Filter 
through a dry filter-paper and wash the precipitate with alcohol. vem 


, 
{ 
F 





MAGNESIUM 99 


4, If the previous manipulation has been faulty, the precipitate obtained 
with perchloric acid-may consist of ammonium perchlorate, and it is advisable, 
therefore, to confirm the potassium test. Dissolve the precipitate on the filter 
with as little hot water as possible, using not over 10 cc. at the most. Add 
the water in small portions around the top of the filter paper and pass the first 
filtrate through the filter a second time. Add 8 to 5 ce. of sodium nitrite 
reagent and 2 to 3 cc. of 6 N acetic acid and boil gently for about five minutes. 
The boiling with nitrous acid serves to decompose any ammonium salt present; 
it will cause the decomposition of as much as 30 mgs. of ammonium ions. 


NH,++NO.-— N, t +2H.0. 


Cool the solution and add a little sodium cobaltic nitrite reagent. A yellow 
precipitate of K,NaCo(NO2). will form within ten minutes if as much as 0.3 
m. of potassium is present. 

5. Test for Sodium. Pour the alcoholic filtrate from the perchlorate pre- 
cipitate into a small Erlenmeyer flask, place the flask in cold water atid saturate 
with dry hydrogen chloride gas. This gas may be prepared by dropping con- 
centrated sulfuric acid into a flask containing common salt and concentrated 
hydrochloric acid and passing the escaping gas through concentrated sulfuric 
acid. In the presence of alcohol the hydrogen chloride will precipitate a 
little as 1 mg. of sodium ions as sodium chloride. Filter off the precipitated 
sodium chloride and wash it with a little alcohol. Dissolve it in a very little 
water, evaporate the solution to dryness, dissolve in 1 cc. of water and add 
twice as much potassium pyroantimonate reagent. A crystalline precipitate 


of sodium pyroantimonate can be obtained with as little as 1 mg. of sodium ions 


if the solution is allowed to stand overnight. Traces of many other elements, 
such as the alkaline earths, magnesium and aluminium also give precipitates 
with this reagent, but such precipitates are flocculent. 

The final precipitates of potassium and sodium should always be submitted 
to the flame test in all cases of doubtful precipitation. 


TaBie II.—ANAtysis oF Group V. Mernop B. 





NH.+ Testthe|_. 
Sri Nets Filtrate from Group IV, Method A, may 


et Method A. contain Mg++, Kt, Nat, NHy+. Concentrate to 10 cc., 
p. 97. ’| add 15 cc. (NH4)2COs and 15 cc. CxHyOH. (1) 





Precipitate:] Filtrate: K+, Nat, NH,+. Remove sul- 
MgCoOs3:(NH,4)s fate by a BaCh, remove Bat+t with 
CO;:4H:0. Dis-| (NH,),CO;. Expel NH, salts and 
solve in 6-normal add HCIO,. (3) 

H»SO, and add an 


equal volume of ae 
C.H;OH. Filter, | Precipitate: KC1O,, | Filtrate, Nat. Test 


add NH,OH and| Examine as in| for Na as in 
NasH PO; to pre- Method A, Dp. 98. Method A, p. 98. 
cipitate Mg 














a Law 


Na ee ee ee 








100 REACTIONS OF THE METALS 


PROCEDURE* 


1. Concentrate the solution to-a volume of about 10 ec. and filter if necessary. 
If ammonium salts are deposited it is best to remove them first as in Method A. 
Add to the concentrated solution 15 cc. of ammonium carbonate reagent and 
15 cc. of alcohol, stir well, and allow the mixture to stand half an hour or longer, 
Filter off the precipitate of MgCO;-(NH,)2.CO;-4H.0. 

2. To confirm the magnesium test, dissolve the precipitate, which may 
contain alkaline-earth carbonate, in a little 6 N sulfuric acid and add an — 
equal volume of alcohol. Any barium, strontium or calcium will be converted 
. into insoluble sulfate by this treatment. Filter the solution if necessary and 
test for magnesium with sodium phosphate solution in the usual way. 

3. Test for Potassium and Sodium. Evaporate the filtrate from the precipi- 
tate obtained by treatment with ammonium carbonate, expel ammonium salts 
and if a sulfate is present, heat to boiling and add hot barium chloride solution — 
until there is no further precipitation. To the filtrate add (NH4)sCO; solution, 
heat to boiling and filter off the BaCO;. Evaporate the filtrate to dryness in a 
porcelain dish and heat the residue until no more white fumes are evolved. 
Cool completely, add 5 cc. of water, washing down the sides of the dish with it, 
filter off the carbonaceous residue, evaporate to dryness and again heat to make 
sure that all the ammonium salts are expelled. Cool, add HCIO,, and treat as 
directed in Method A, p. 98. If no sulfate is present it is unnecessary to add 
BaCl, and (NH,)2COs;, otherwise the procedure is the same. 


The separation of the potassium and sodium can be accomplished 
by treatment with chloroplatinic acid instead of perchloric acid, but 
such a method is not suitable in ordinary qualitative analysis on account 
of the expense involved. The treatment is practically the same except 
that instead of evaporating till fumes of perchloric acid are evolved, 
the solution is evaporated just to dryness on the water-bath. 


*Cf. Scuarrcorr, Ann. Phys., 194 (1858), 482; Goocu and Eppy, Z. anorg. 
Chem. (1908), 427; A. A. Noyes, A Course of Instruction in the Qualitative 
Analysis of Inorganic Substances (1914). 





Y 


GROUP IV. ALKALINE EARTHS 


CALCIUM, STRONTIUM, BARIUM 


GENERAL CHARACTERISTIC REACTIONS 


The metals of the alkaline-earth group are bivalent and heavier 
than water, which they decompose slowly at ordinary temperatures, 
with evolution of hydrogen and formation of difficultly soluble hydrox- 
ides of strongly alkaline reaction (cf. p. 41). The salts are mostly 
colorless and very slightly soluble in water. The halogen compounds, 
nitrates, nitrites, and acetates are soluble in water. The carbonates are 
insoluble in water and are decomposed on ignition into carbon dioxide, 
and white, infusible, strongly luminous metallic oxide: 


CaCO3 = Ca0+C02}. 


Strontium carbonate is less readily decomposed than calcium carbonate, 
and barium carbonate loses its carbon dioxide only when heated to a 
white heat; its oxide is not very luminous. 

The sulfates and oxalates are very difficultly soluble. The 
sulfate of barium is the most insoluble sulfate and calcium sulfate the 
most soluble; of the oxalates, the calcium salt is the most insoluble 
(ef. p. 21). The solubility of the strontium salt is always midway 
between that of the corresponding calcium and barium salt, for the 
atomic weight of strontium, of which the solubility is a function, lies 
midway between the atomic weights of barium and calcium. The halo- 


gen salts are volatile and impart a characteristic color to the flame. 


The metals of this group form oxides of the general type RO, and 
peroxides corresponding to the formula ROz. The latter, on treatment 
with acids, give hydrogen peroxide and salts corresponding to the oxide 


RO: 
ROe+2HCl = H202+RCle. 


Magnesium is more closely related to this group than to the alkali 
metals. It can be precipitated with this group if the group precipitant, 
ammonium carbonate, is added to the concentrated solution together 
with an equal volume of alcohol. 

101 


102 REACTIONS OF THE METALS | 


CALCIUM, Ca. At. Wt. 40.07 
Sp. Gr. 1.58. M. Pt. 810°C. 


Occurrence.—Calcium is widely distributed in nature. It is found 
in enormous deposits in all stratified formations as carbonate (limestone, 
marble, chalk), often rich in petrification. The carbonate, CaCOs, 
is dimorphous, crystallizing in rhombohedrons as calcite and in the 
orthorhombic system as aragonite. Calcium also occurs in large 
masses as sulphate, partly as monoclinic crystallizing gypsum, 
CaSO.4-2H20, and partly as anhydrite, CaSO4, which crystallizes in the 
orthorhombic system. Calcium also occurs as fluorite, CaFe, which 
crystallizes in the isometric system, with perfect octahedral cleavage; 


Cl 
as apatite, 3Cag(PO.)2: Cat , which belongs to the hexagonal system; 
KF . 


and, finally, in innumerable silicates, such as the monoclinic wollaston- 
ite, CaSiO3, and the triclinic anorthite, CaAlsSis0s. The calcium 
minerals are the principal representatives of several important 
isomorphous groups: 


Calcite Group, Rhombohedral. Aragonite Group, Orthorhombie. 


Calcite, CaCOsz Aragonite, CaCOz 
Magnesite, MgCOs Strontianite,  SrCOz3 

: Ca Witherite, BaCOsg 
pepomite, Mg as Cerussite, PbCOg3 
Siderite, FeCO3 


Smithsonite, ZnCOz3 
Rhodochrosite, MnCOg 


Anhydrite Group, Orthorhombic. Apatite Group, Hexagonal. 


Anhydrite, CaSO4 Apatite, 3Caz(PO4)2+Ca(ClF) 
Celestite, SrSO4 Pyromorphite, 3Pb3(PO4)2+PbCle 
Barite, BaSO, Mimetite, 3Pb3(AsO4)2+PbCle 
Anglesite, © PbSOs Vanadinite, 3Pb3(VO4)2+PbCle 


REACTIONS IN THE WET WAY 


1. Ammonia, in case it is free from carbonate, produces. no pre- 
cipitate with calcium salts; on standing in the air, however, carbonic 
acid is absorbed and a turbidity of calcium carbonate results. 

2. Ammonium Carbonate, or any other soluble, normal carbonate, 
precipitates white calcium carbonate; the precipitate is voluminous 





CALCIUM | 103 


and flocculent when it first forms, but soon becomes crystalline, par- 
ticularly when in contact with boiling water. 


Cat++CO37* — CaCO3. (1) 


‘The precipitate is noticeably soluble in an aqueous solution of the 


ammonium salt of any strong acid: 
CaCO3+2NH4t — Catt++2NH3 T 41,0 +00: 1 a (IT) 


When ammonium carbonate is the precipitant, equation (II) is essen- 
tially the cpposite to equation (I) and the mass action law shows how 
the reaction can be made to go in either direction. An excess of am- 
monium carbonate will favor the progress of equation (1) and boiling 
with a large quantity of an ammonium salt such as ammonium chloride 
will cause equation (II) to go to completion. 


Ammonium carbonate is an unstable substance (cf. p. 51). Commercial 
ammonium carbonate, often called ammonium sesquicarbonate, is a mixture 
of approximately equivalent quantities of ammonium bicarbonate, NH,HCOs, 
and ammonium carbamate, NH.CO.NH2; the latter salt corresponds to normal 
ammonium carbonate less one molecule of water. Calcium carbamate is quite 
soluble in water, but by contact with water at 60° it becomes changed to insolu- 
ble calcium carbonate. The ammonium carbonate reagent is prepared with 
6 N ammonia instead of water; this prevents the hydrolysis of the salt and 
changes the bicarbonate to ammonium carbonate. 


Calcium bicarbonate is soluble in water, so that any acid which is 
dissociated to a greater extent than HCOs3 will exert a solvent effect 
upon calcium carbonate. 


CaCO; +H Fi Cat Fi. HCO3. 


Thus acetic acid dissolves calcium carbonate readily. Boiling the solu- 
tion favors the progress of the reaction, as the HCO37 is also in equi- 
librium with H* and H2CO3 and the latter with H2O and COs: 


HCO;"+H* — HeCO3 — H20+COd2, 


and the carbon dioxide can be expelled completely at 100° (ef. p. 15). 

The precipitation of calcium carbonate from boiling, dilute solutions 
containing ammonium salts is always more or less incomplete, for 
the reasons that have just been given, but from cold, concentrated 
solutions containing considerable alcohol the precipitation is practically 
complete. 

3. Ammonium Oxalate produces in neutral or alkaline solutions 
a precipitate of calcium oxalate, which when formed from cold solutions 


104 i: REACTIONS OF THE METALS 


is composed of extremely fine crystals, hard to filter, while from hot 
solutions larger crystals are formed: 


C2047 +Catt = CaC20,. 


Calcium oxalate is practically insoluble in water and aectin acid, but 
dissolves readily in mineral acids: 


CaC.04+2Ht 2 Catt++HoCs0u. 


Calcium oxalate, unlike calcium carbonate, does not dissolve in 
acetic acid for the following reasons: (1) The solubility product of 
calcium oxalate is about 3.8x10-® while that of calcium carbonate is 
about 1.7X10-® (p. 21) which shows that the oxalate is somewhat 
less soluble in water than the carbonate. (2) The ionization constant 
of HCO3° is 0.0638 while that of HC2047 is 0.045 and that of acetic acid 
is 0.0418; this reason is sufficient to explain why acetic acid has little 
solvent effect upon calcium oxalate. To dissolve calcium oxalate readily 
it is necessary to use an acid strong enough to repress the ionization of 
the first hydrogen of oxalic acid, for which the ionization constant is 
0.038. (38) The progress of the reaction cannot be aided, as in the case 
of the carbonate, by the loss of a volatile constituent. 

Ammonia precipitates from such a solution calcium oxalate again, 
the excess of hydrogen ions, as well as, the oxalic acid which was formed, 
being neutralized. 

Calcium oxalate on being boiled with sodium carbonate solution 
is easily changed to carbonate: 


CaC204+C0O37 @ CaCO3+ C2047. 


This reaction takes place in the direction left to right when an excess 
of COs” ions are present in spite of the fact that calcium oxalate is 
somewhat less soluble than calcium carbonate. This is in accordance 
with the mass action law (p. 13). An excess of C2047 ions will make 
the reaction take place in the direction right to left. | 

4, Sulfuric Acid produces a precipitate only in concentrated 
solutions: 


CaCle +He2SO4 = 2HCl+CaSO.. 


One hundred cc. water disolves 0.214 gm. CaSO4-2H2O at 40°, 
but much less than 1 mg. will dissolve in the same quantity of alcohol. 
In the presence of a slight excess of sulfate ions, calcium sulfate is less 
soluble, but the behavior of calcium sulfate toward hydrogen ions is 
similar to that of calcium carbonate. Calcium acid sulfate is much 
more soluble than normal calcium sulfate, and any acid which is capa- 





sa eee 
; 7 - 


CALCIUM , 106 


ble of repressing the ionization of HSO,7 will exert a solvent effect 
upon CaSO4. Sulfuric acid itself, or any other strong acid, can exert 
this solvent effect. 

~ Calcium sulfate is also soluble in concentrated ammonium sulfate, 


owing to the formation of a complex anion: 


CaSO4+(NH4)2804 => (NH4)2[Ca(SO4)o]. 


5. Calcium Sulfate solution produces no precipitation with calcium 
salts. (Note difference from strontium and barium.) 

6. Sodium Phosphate (NazHPOs) produces in teaibcal solutions 
a white, flocculent precipitate of secondary calcium phosphate: 


Catt++HPO,- — CaHPOx,. 


If ammonia is added to the solution at the same time, tertiary 
calcium phosphate will be precipitated: 


HPO, +OH — He0+P0.4- 
Cat t++2P04* => Cazg(POx)o. 


Both of these phosphates of calcium are dissolved by hydrogen ions. 
A glance at the table on p. 10 shows that even acetic acid will 
repress the ionization of H2PO4" to an extent such that it will exert a 
slight solvent action upon CaHPO, and upon Cag(POxs)o. Acetic 
acid in the presence of a soluble acetate, however, will not dissolve 
Caz(PO4)2 to any extent. From a solution obtained by dissolving cal- 
cium phosphate in acid, ammonia always precipitates the tertiary salt. 

7. Chromates of the Alkalies do not precipitate calcium salts 
from dilute solution. (Note difference from barium and strontium.) 

8. Absolute Alcohol, or a mixture of equal parts of absolute alcohol 
and ether, dissolves both the nitrate and chloride of calcium. 

All deliquescent salts, with the exception of potassium carbonate, 
dissolve in absolute alcohol. All other salts are, in general, insoluble, 
or very difficultly soluble, in absolute alcohol. An exception to this 
rule is found in the case of mercuric chloride, which is not deliquescent, 
and is much more readily soluble in alcohol than it is in water. 

9. Water decomposes the carbide, phosphide, and nitride of calcium 
at the ordinary temperature, as follows: 

(a) The carbide: 

CaC2+2HOH = Ca(OH) 2-+CoHp 1. 


Acetylene 
Acetylene is evolved by the reaction, a gas with a peculiar smell.* 


* Pure acetylene is odorless. Almost all calcium carbide contains a little ecal- 
cium phosphide, which evolves phosphine on treatment with water. 





106 REACTIONS OF THE METALS 


If this gas is conducted into an ammoniacal copper solution, it rapidly 

produces a red precipitate of copper acetylide. The latter compound 

is harmless while it is moist; but in the dry state it can be readily 

exploded by a blow, by rubbing, or by simply warming. 
(b) The phosphide: 


- Ca3P2+6HOH =3Ca(OH)2+2PH3 1. 


The garlic-smelling phosphine gas which is evolved is spontaneously 
combustible, because it always contains a small amount of the spon- 
taneously combustible liquid, phosphuretted hydrogen (P2H4). 

(c) The nitride: 


CasN2+6HOH =3Ca(OH)2+2NH Tf 3. 


REACTIONS IN THE DRY WAY 


Calcium compounds, on being heated with sodium carbonate 
before the blowpipe, are changed to the white infusible oxide, which 
glows brightly when hot. 

The volatile calcium compounds color the non-luminous gas flame 
brick-red. 

Flame Spectrum.—Orange-yellow double line (620.3 uy, 618.2 pp) 
and a yellowish-green one (554.4 wy, 551.8 wu); both of these lines 
belong to calcium oxide. If calcium chloride wet with hydrochloric acid 
is placed in the flame, a number of other lines are seen; in the orange- 
yellow 646.6 uu, 606.9 uu, 604.5 yu and 593.4 wy, in the yellow 581.7 uy 
and 572.0 uu, in the violet, usually very hard to see, 422.7 yy (see 
spectroscopic chart, Frontispiece), 


STRONTIUM, Sr. At. Wt. 87.63 
Sp. Gr. 2.5. M. Pt.>Ca. and<Ba. 


Occurrence.—Strontium occurs quite commonly with calcium, 
but usually in much smaller amounts. There are only a few true 
strontium minerals. The most important of these are: Strontianite, 
SrCOz, orthorhombic, isomorphous with aragonite; and celestite, 
_ SrSO4, orthorhombic, isomorphous with barite. 


REACTIONS IN THE WET WAY 


1. Ammonia: same as with calcium. 

2. Ammonium Carbonate: same as with calcium. 

3. Ammonium Oxalate: same as with calcium; but the stron- 
tium oxalate is somewhat soluble in acetic acid. 


STRONTIUM 107 


4, Dilute Sulfuric Acid produces a white precipitate of -stron- 
tium sulfate: 
SrCle+-H2SO4 = 2HCl+SrSOx. 


Strontium sulfate is much less soluble in water than calcium . 
sulfate (6900 parts of water at ordinary temperatures dissolve 1 part 


F $rSO.z), but much more soluble than barium sulfate. It is soluble in 





boiling hydrochloric acid, and insoluble in ammonium sulfate. By 
boiling with a solution of ammonium or alkali carbonate solution, the 
strontium sulfate is changed to carbonate: 


$rS0O.1+COs  — SrCO3+S0.". 


5. Calcium Sulfate solution produces in neutral or weakly acid 
solutions, after some time, a precipitate of strontium sulfate: 


Srt *+-CaSO4 =SrSO4+Cat?. 


-6. Chromates of the Alkalies produce in dilute solutions no pre- 
cipitate (thus differing from barium); but from concentrated solutions 
strontium chromate is precipitated. It is much less soluble in alcohol; 
100 ce. of alcohol, 53 per cent by volume, will dissolve 0.002 gm. SrCrO4 
and 0.088 gm. of CaCrO4 at room temperature; 100 cc. of 29 per cent 
alcohol will dissolve 0.1382 gm. SrCrO4 and 1.22 gms. CaCrO4. The 
precipitate is quite soluble in acetic acid. 

7. Absolute Alcohol. The nitrate is not deliquescent, and does not 
dissolve in absolute alcohol. Strontium chloride is slightly deliquescent; 
the anhydrous salt dissolves scarcely at all in absolute alcohol; but, 
on the other hand, 1 gm. of SrClz-+6H2O dissolves in 116.4 gms. of cold 
alcohol and 262 gms. of boiling absolute alcohol. 100 cc. of 66 per 
cent alcohol dissolves about 50 gms. of SrCle-6H20. 


REACTIONS IN THE DRY WAY 


Heated on charcoal before the blowpipe the strontium compounds 
behave similarly to the calcium compounds. 

The volatile strontium salts color the non-luminous gas flame 
carmine red. | 

Flame Spectrum.—A number of lines in the red and orange yellow 
and one in the blue. No bands in the green. Red 686.3 uy, 674.7 up, 
662.8 wu, 649.9 wu; orange-yellow 646.5 wu, 635.1 uu, 606.0 uu; blue 
460.7 uy. 


108 REACTIONS OF THE METALS 


BARIUM, Ba. At. Wt. 137.37 
Sp. Gr. about 4.0. M. Pt. 850° (?) C. 


Occurrence.—Like strontium, barium is almost always found 
associated with calcium, but only in small amounts. The most im- 
portant barium minerals are: Witherite, BaCO3, orthorhombic, 
ismorphous with aragonite; barite, or heavy spar, BaSO4, ortho- 
rhombic, isomorphous with anhydrite; and the hydrous barium 
aluminium silicate, harmotome, BaAleH2Sis0i5+4H20. Harmotome ~ 
crystallizes in the monoclinic system, and belongs to the class of zeolites. 


REACTIONS IN THE WET WAY 


1. Ammonia and Ammonium Carbonate: same as with calcium 
and strontium. 

2. Ammonium Oxalate: same as with calcium and strontium, 
except that the barium oxalate formed is more soluble in water (1 gm. 
dissolves in 2.6 liters of cold water), and is readily dissolved by hot 
dilute acetic acid. 

3. Phosphates of the Alkalies: same as with calcium. 

4. Chromates of the Alkalies produce in neutral solutions of barium 
salts a yellow precipitate of barium chromate (thus differing from cal- 
cium and strontium), 


Bat to CrO4” = BaCrOx,. 


The table on p. 10 shows that HCrO,- is ionized only to between 0.1 
and 0.2 per cent in 0.1 N solution. In the presence of a stronger acid, its 
ionization becomes much Iéss, and the equilibrium 


Ht+cCrO,- —@ HCrO; and 2HCrO,- — H.O+Cr.0;= 


progresses farther in the direction left to right than in pure water. The equilib- 
rium, however, also depends upon the concentration of CrO,- in the solution. 
In the case of strontium chromate, enough CrO, ions are present in the satu- 
rated aqueous solution so that the formation of the HCrO, takes place to a 
considerable extent when a little acetic acid is added to the solution, and it is 
easy in this way to prevent the precipitation of as much as 0.5 gm. of strontium 
ions; barium chromate, on the other hand, is so much more insoluble that it 
takes considerable acetic acid to have an appreciable effect upon its solubility. 
Thus a small quantity of barium can be separated from quite a large quantity 
of strontium by adding CrO¢ and a suitable quantity of acetic acid. 

If a more highly ionized acid is present, such as hydrochloric acid, the barium 
will not be precipitated as chromate, but by adding sodium acetate (cf. p. 46) 
the concentration of the hydrogen ions can be reduced sufficiently to precipi- 
tate even a small quantity of barium as chromate. 








BARIUM 109 


5. Dilute Sulfuric Acid produces, in even the most dilute solutions, 
a precipitate of barium sulfate: 


Bat *++S04- — BaSO,. 


According to the table on'p. 21, 1 liter of pure water dissolves about 
2.5 mgs. of barium sulfate. In a slight excess of sulfuric acid it is 
much less soluble by virtue of the common ion effect (p. 45). As with 
strontium and barium sulfates, the presence of an excess of hydrogen 


- ions has a solvent action due to the formation of acid sulfate. This 


effect is appreciable with acids such as hydrochloric or nitric acid, but 
the solubility of barium sulfate is so slight that it requires treatment 
with hot, concentrated sulfuric acid in order to get any considerable 
quantity of barium sulfate into solution, and dilution with water causes 
reprecipitation of barium sulfate: 


BaSO, +H2S804 2 Ba(HS04)s. 


Barium sulfate is partially converted, as a result of the mass action 
effect, into more soluble barium carbonate by boiling with a concen- 
trated solution of sodium carbonate: 


BaSO.+ NaeCO3 —@ BaCO3 +Na2SQO4. 


To make this decomposition quantitative, the barium sulfate 
must be boiled with the sodium carbonate solution, filtered, treated 
with a new portion of sodium carbonate solution, and the process re- 
peated until the filtrate no longer gives a test for sulfate. The more 
concentrated the sodium carbonate solution is, the more complete 
will be the decomposition. The highest degree of concentration will be 
reached by fusion of the barium sulfate with anhydrous sodium carbonate. 


Consequently, to obtain a solution of barium ions from insoluble barium 
sulfate, it is best to proceed as follows: Mix the solid with four to six times as 
much calcined sodium carbonate and fuse the mixture in a platinum crucible. 
Cool, boil the residue with a little water until thoroughly disintegrated, and 
filter. Wash the residue with hot, normal sodium carbonate solution, until the 
filtrate gives no test for sulfate ions, and then with a little water. Dissolve 
the residue of barium carbonate in dilute hydrochloric, nitric or acetic acid. 

If the product of the fusion were treated with considerable water, or if the 
residue were washed at once with considerable hot water, the dissolved sodium 
sulfate would react with the insoluble barium carbonate to form less soluble 
barium sulfate. This is prevented, in accordance with the mass-action prin- 
ciple; by keeping the concentration of the sodium carbonate sufficiently large. 


6. Fluosilicic Acid produces a white, crystalline precipitate of 
barium fluosilicate: 


H2SiF.+BaCle =2HCl+ BaSiFs. 


110 REACTIONS OF THE METALS 


In order to effect complete precipitation, the solution must stand some — 
time. Barium fluosilicate is difficultly soluble in water and dilute 
acids, and insoluble in alcohol. 

7. Absolute Alcohol dissolves neither the nitrate nor the chloride; 
neither of these salts is deliquescent: 100 cc. 66 per cent alcohol (by 
volume) dissolve 3.3 gms. of BaCle-2H20. | 

8. Concentrated Hydrochloric Acid and Nitric Acid will precipitate 
from fairly-concentrated barium solutions the chloride and nitrate 
respectively. | 


REACTIONS IN THE DRY WAY 


Heated with sodium carbonate on charcoal before the blowpipe, the 
barium compounds, unlike those of calcium and strontium, do not give 
a brightly luminous mass, because the barium carbonate formed is not 
- decomposed at this temperature into the infusible oxide and carbon 
dioxide, but merely sinters together somewhat. Volatile barium salts — 
color the non-luminous gas flame yellowish-green. ,\ The sulfate is 
only slightly volatile in the hottest flame, and in the ordinary gas flame 
it shows scarcely any coloration. In order to obtain the coloration, 
it is best to change the sulfate into chloride, by reducing a small 
particle on a platinum wire in the upper reducing flame to sulfide and 
adding a little hydrochloric acid by means of a capillary tube. The 
wire on then being brought into the flame will give the characteristic 
flame coloration. 

Flame Spectrum.—A number of deep green lines, weaker lines in 
the orange-yellow part of the spectrum, and one blue line. 

Orange-yellow (654.0 uy, 629.8 pu), (624.0 uu, 617.9 pu, 610.9 wp, 
603.2 wu). Yellow 528.5, yellow triple line (576.9 uy, 572 wy, 564.8 up). 
Green 553.5 wy, 534.7 wy, 524.3 wy, 513.7 wy, 500.0 wu. Blue, 484.7 pp. 


Separation of Calcium, Strontium, and Barium 


In the course of a systematic analysis these three metals are always 
obtained in the form of their insoluble carbonates, either by precipita- 
tion with ammonium carbonate from fairly dilute solution in the presence. 
of ammonium chloride (cf. p. 103) or by fusion of the sulfates with 
sodium carbonate. 

A number of excellent methods have been proposed for the analysis 
of the alkaline earth group. The characteristic reactions of barium, 
strontium and calcium are so similar that difficulties are likely to arise 
in any scheme whenever a small quantity of one of these elements is 
present together with a large quantity of another; thus a precipitate 





STE a ee 





Ne ee, le ee be 
Pai ee oe 


TAS 


BARIUM 111 


caused by the presence of much strontium may be confused with one. 
produced by a little barium. Two methods of analyzing the group will 
be described. 


TasBLeE III.—Anatysis or Groupe IV. Mernop A. 





Solution may contain: Bat+, Srt+, Ca++, Mg++, K+, Nat, NH«t. Add 
NH,OH and (NH;,)2CO;. Examine filtrate for Mg, K, Na according to Table I, p. 97, or 
Table II, p. 99. (1). Dissolve carbonates of Ba, Sr and Ca, in 2-normal HNOs3. 
Evaporate to dryness. Dissolve a part of the residue in a little water and add CaSO, 
solution (2); (a) no precipitate is formed. Ca is present (3); (b) a precipitate forms 
slowly, Sr is present and possibly Ca (4); (c) a precipitate is formed at once, Ba is 
present and possibly Sr and Ca (5). Treat the remainder of the dry residue with 
C,H;OH (6). 





Residue: Sr(NOsz)2, Ba(NOs)2, Heat with solid | Solution: Ca(NOs3)2. Evap- 
s NH.,Cl (8). Treat cold residue with C2H;OH. orate to dryness and test 
in the flame.  Brick-red 
coloration shows Ca. (7) 





Residue: BaClh. Test in | Solution: SrCl,. Hvaporate 
the flame. Dissolve in| to dryness and test in 
water, heat to boiling and| flame. Carmine-red 
add HC.H;0.and K2CrO,,| coloration shows Sr. (9) 
yellow precipitate of ‘ 

BaCrO, shows Ba. (9) 











PROCEDURE ’ 


1. Concentrate the filtrate from Group III to small volume in a porcelain 
dish, add 6N HCl to acid reaction and filter if the solution is not perfectly 
clear. If sulfur runs through the filter paper, make a pulp by shaking small 
pieces of filter paper in a bottle with water, add some of this pulp to the turbid 
solution, stir or shake, and filter. Wash the residue till the volume of the 
filtrate is about 50 cc., heat this solution to boiling, add NH,OH and 
(NH,).CO;. Filter and examine the filtrate for Mg*t, Kt and Nat. The 
precipitate may contain BaCO;, SrCO; and CaCQs. 

2. Dissolve the carbonates in dilute nitric acid and evaporate the solution 


- in a small porcelain crucible just to dryness, by heating with a small flame kept 


in constant motion and heating very carefully until all the moisture and nitric 
acid have been expelled. Take care not to overheat the residue, as this will 
cause the conversion of the nitrates to oxides. Dissolve a small portion of the 
residue in as little water as possible and add 5 cc. of calcium sulfate solution. 

3. If no precipitate is formed, only calcium (and magnesium) can be present. 

4. If a precipitate forms only after standing some time, barium is absent, 
strontium is present and calcium may be present. 

5. If a precipitate forms immediately, barium is present and possibly the other 
members of the group as well. This preliminary test with calcium sulfate 
solution is very useful. It must be remembered in making this test that 5 ce. 
of saturated calcium sulfate solution contain only 3 mg. of calcium ions so that 
a large precipitate cannot be obtained with this reagent even when much barium 
is present; 10 mgs. of barium will give the maximum possible precipitate with 
5 ec. of calcium sulfate reagent. 


112 REACTIONS OF THE METALS 


6. Treat the.remaining residue, which must be perfectly dry, with a little 
absolute alcohol, stir with a glass rod and pour the alcohol through a filter which 
has been wet with absolute alcohol, catching the filtrate in a small porcelain 
crucible. If any calcium was present in the original residue of nitrates, some 
of it will have been dissolved by the alcohol. 

7. Evaporate the alcoholic solution to dryness, wipe out the crucible with 
a piece of filter paper containing no calcium, fold the paper and fasten it to a 
platinum wire. Burn the paper, moisten the ash with hydrochloric acid and 
hold it in the non-luminous flame. The presence of calcium is shown by a 
brick-red flame coloration. 

8. If calcium is found to be present, repeat the treatment of the original 
nitrate residue with absolute alcohol several times to dissolve out all the cal- 
cium nitrate. Then mix the remaining residue with an excess of ammonium — 
chloride and heat until no more ammonium fumes are evolved. By this 
operation the nitrates are converted to chlorides. The nitrogen, which has a 
negative polarity of three in ammonium chloride (having four negative valences 
and one positive) is oxidized by the quinquevalent, positively-charged nitrogen 
of the nitrate and the product of this oxidation and reduction is nitrogen. At 
the same time the negatively-charged chlorine is oxidized by the nitrogen of 
the nitrate, forming free chlorine and nitric oxide, the latter escaping partly 
as such and partly oxidizing the nitrogen of the ammonium. The oxidation 
of the ammonium chloride requires four positive charges of electricity; the 
nitrogen of the strontium nitrate loses five positive charges if reduced to nitro- 
gen and three if reduced to NO. Assuming a complete reduction to nitrogen, 
the equation is 


28r(NO;).+6NH,Cl =28rCl,+5N, 7 +Cl Tt +12H,0; 
assuming the nitrate reduced only to nitric oxide, NO. 
Sr(NOs)2-+2N HCl =SrCl,+4H20+2NO 1 +N. 7. 


It is also easy to change the nitrate to chloride by evaporation with con- 
centrated hydrochloride (cf. p. 57). 

9. Treat the residue of chlorides with a little 66 per cent alcohol (by vol- 
ume) exactly as described above and test the alcohol solution for strontium by 
the flame test; a carmine-red coloration shows the presence of strontium. Wash — 
any residual chloride with 80 per cent alcohol to dissolve out all the strontium 


chloride and test the residue likewise in the flame for barium; a yellowish-green _ 


coloration shows barium. If there is sufficient residue, dissolve it in a little water 
and a little acetic acid, heat to boiling and test for barium with potassium 
chromate solution; a yellow precipitate at this stage of the analysis is es 
proof of the peice of barium. 


The above cpthad of analysis is based on the assumption that the 
alkaline-earth carbonates were formed under conditions such that only 
barium, strontium and calcium are present in the carbonate residue or 
precipitate. If the ammonium carbonate is added under the conditions 
described on p. 100, the magnesium will be precipitated as carbonate, 
together with calcium, strontium and barium. The following method 





BARIUM 113 


of analyzing this group has been tested thoroughly by A. A. Noyes* 
and his students and permits the identification of as little as 1 mg. of 
any constituent in the presence of 500 mgs. of any other member of the 
group. 


Taste ITV.—ANALysis oF Group IV. Mernuop B. 





Solution may contain: Bat+, Srt+, Mg++, K+, Nat, NHit+. Concentrate to 
10 cc.; add 15 cc. (NH1)2COs reagent, or more if necessary, and an equal volume of 
C.H,OH. Stir, let stand 30 minutes and filter. Test filtrate for Na+ and K* accord- 
ing to Table I, p. 97. Dissolve the precipitate, which may contain BaCO;, SrCOs, 
CaCO; and MgCO3-(NH3)2CO3,-4H20, in 6-normal HC2H;0. add NH,C2H;O0. and 
K2CrO,. (1) . 





Precipitate:| Filtrate: Sr++, Cat+, Mg++. Add NH,OH. Dilute to 65 cc. 
BaCrO,. Dissolve} and add 50 cc. C:H;OH. Shake the solution with filter paper 
in HCl. LEvap-| pulpand filter. (8). . : 
orate to dryness. 


- Test residue in Li 
flame, treat with Precipitate: 





Filtrate: Cat +, Mg++. Dilute to 150 ce. 


4 
? , 
y 
: 


3 CC. of 6-normal SrCrOx. Pour a 
Minriwre os 
(NH,)2COz; and 
K2CrO, through 
the filter and reject 
this filtrate. Wash 
with cold water 
and dissolve the 
residue of SrCO3 
inatlitttle 
HC.H;0.. Test 
with CaSQx,. Slight 


HC2H;02, 20 cc. 
of 3-normal 
N H. 4CoH AOD and 
15 cc. of water. 
Heat to boiling 


and test’ with 


K2Cr2O,4_ solution. 
Yellow precip- 
itate is BaCrQy,. 


— (2) 


precipitate, shows 


Sris present. (4) 


heat to boiling andadd (NH,4)2C20,. (5) 





Precipitate: 


CaC.0,4. Dissolve 
in 5 ce. of 6- 
normal HySO. and 
add 10 cc. of 
C.H;OH. White 
precipitate 
is CaSO... (5) 


Filtrate: Mgtt. 


Add NH,OH and 
N. aoH P O14. Dis- 
solve in 2-normal 
HSO,, add 20 ce. 
C.H;OH and. filter 
af necessary. Add 
NH,OH. Precipi- 
tateis MgN HyPQO,. 
(6) 


lk et i ee a i ae 
= Yo 














PROCEDURE 


1. Dissolve the fairly dry carbonates in hot, 6 N acetic acid and evaporate 
the resulting solution just to dryness, taking care not to overheat the residue. 
Moisten the residue with 3 cc. of 6 N acetic acid and dissolve it in 20 ce. 
of water. Add 20 cc. of 3 N ammonium acetate solution, heat to boiling 
and precipitate the barium as chromate by adding hot 3 N potassium 
chromate solution, drop by drop, until no further precipitation takes place 
and the solution is decidedly yellow in color. Boil gently for about two minutes 
longer and filter through paper capable of holding very fine crystals. 

2. If the-filtrate is not decidedly orange in color, showing the presence of an 
excess of dichromate ions, add a little more chromate to see if the precipitation 
of the barium is complete. Wash the precipitate thoroughly with cold water, 
rejecting all but the first of the washings. To confirm the barium test, dissolve 





“A Course of Instruction in the Qualitative Analysis of Inorganic Substances. 


114 REACTIONS OF THE METALS 


the precipitated barium chromate in a little hydrochloric acid, evaporate the 
solution to dryness and try the flame test (ef. p. 110). If the flame test is incon- 
clusive, it may be that the precipitate contained strontium chromate, due to 
the presence of a large quantity of strontium. In such cases repeat the precipi- 
tation with potassium chromate, using the same quantities of acetic acid and 
ammonium acetate as before. By evaporation to dryness, the chromate is 
_ almost entirely changed to chloride and chromic salt and in the second pre- 
cipitation with chromate in acetic acid solution, no strontium chromate will be 
deposited, because the quantity of strontium now present cannot be large enough 
to come down with the barium a second time. 

3. To test for strontium, make the filtrate from the first precipitation with 
potassium chromate distinctly ammoniacal, dilute to 65 cc. and add 50 ce. of 
alcohol. If necessary add a little more potassium chromate and make sure 
that the solution remains distinctly ammoniacal. Shake the solution with 
some filter paper pulp and then filter. Allow the filter to drain thoroughly, but 
do not wash the precipitate. Place the filtrate aside for the calcium and mag- 
nesium tests. 

4, Pour through the filter, which presumably contains strontium chromate, 
a hot mixture of 10 cc. ammonium carbonate reagent and 5 cc. of potassium 
chromate reagent. This mixture serves to convert strontium and calcium 
chromate to the less soluble carbonates. If barium chromate is present it is 
unaffected by the treatment. Wash the residue with cold water until the wash- 
ings are colorless, which causes the removal of any remaining magnesium. 
Dissolve the carbonate on the filter in a little normal acetic acid and test the 
solution for strontium with calcium sulfate (ef. p. 111). 

5. To test for calcium, dilute the filtrate from the strontium chromate pré- 
cipitate to about 150 cc., heat to boiling and add ammonium oxalate solution, 
in small portions, until no further precipitation takes place. Filter the solution 
while still hot. The calcium is precipitated as oxalate and the magnesium 
remains in solution. To confirm the calcium test, dissolve the precipitated 
calcium oxalate, ora part of it, in 5 cc. of 6 N sulfuric acid and add 10 ee. of 
alcohol. A white precipitate of calcium sulfate shows calcium to be present. 
If the analysis is properly conducted, very little strontium, if any, will be pre- 
cipitated with the calcium as oxalate and practically none of it will dissolve in 
the sulfuric acid. The calcium sulfate test is shown with 0.2 mg. of calcium 
ions if the solution is allowed to stand ten minutes. Not more than 0.3 mg. 
strontium will be present if the above directions are followed properly. Mag- 
nesium does not interfere with the test. 

6. To test for magnesium, make the filtrate from the calcium oxalate precipi- 


*- tate strongly ammoniacal and add a little sodium phosphate reagent (cf. p. 95). 


To confirm the magnesium, dissolve the precipitate in 5 cc. of 2 N sulfuric 
acid, add 20 cc. of alcohol and stir well for two or three minutes. Filter off 
any calcium sulfate that may form, and repeat the precipitation of the i 
- slum as phosphate (cf. p. 98). 


Traces of alkalies and alkaline earths are recognized best by means 
of the spectroscope. At this place, therefore, a brief description of 
spectroscopic analysis will be given. 





BARIUM 115 


Spectroscopic Analysis (Bunsen and Kirchhoff, 1865) 


_ If a ray of white light is passed through a glass prism, not only 
is the direction of the ray changed, but the white light is decomposed 
into colors; it suffers dispersion. It 
will be found that the red rays are 
deflected least, while the violet rays 
are deflected most. The picture ob- 
tained—the spectrum—if projected on 
a screen (Fig. 5), does not show the 
colors sharply separated, but merging 
into one another. Such a spectrum 
“is called a continuous, or uninter- 





rupted, spectrum. LHvery glowing solid Rea 
or liquid body emits white light; the Yellow 
spectra obtained in all such cases will Violet 


be continuous ones. Glowing vapors 
and gases behave quite differently. 
They do not emit white light, but Fia. 5. 
light composed of rays of definite 
wave lengths, which are characteristic for each gas and for each 
vapor. The light emitted from glowing vapors or gases, when 
decomposed by the prism, yields on the screen a discontinuous or 








interrupted spectrum. If the light is passed through a fine slit before 
reaching the prism, the spectrum will be found to consist of a greater 
or less number of colored lines which always appear in the same place 
with any given substance, provided the prism or its position is not 
changed. In order to determine the exact position of these lines, we 


116 REACTIONS OF THE METALS 


make use of the spectroscope of Bunsen and Kirchhoff (Fig. 6). Fig. 7 
shows a cross-section of the apparatus. 
s : 


O 





P 





P 


Sk 





























Fria. 7. 











The substance to be examined is placed in the loop of a platinum 
wire and introduced at A into the non-luminous gas flame, by means 
of which it is volatilized. The rays of light pass through the slit 
into the tube Sp, reach the prism, by which the rays are refracted 


BARIUM | 117 


into the telescope C (the collimator) and are observed at D. Upon 
a glass plate at the end of the tube Sk is a transparent scale, which is 
illuminated by a small flame at B. This tube is so tnelined toward 
the face aa of the prism that the rays of light from this tube are totally 
reflected into the tube C, and reach the eye of the observer; thus the 
rays from the substance appear at a certain position on the scale. 
As, however, the position of the lines depends upon the dispersive 
power of the prism, and upon its angle of refraction, it is clear that the 
position of the lines will be somewhat different in different spectroscopes. 
_ As every ray has a definite wave length, it is better to give the wave 
__ lengths of the rays which appear, rather than their position on the scale. 
Wave lengths are expressed in millionths of a millimeter, called 
5 





micromillimeters, and designated as up. 
The following values for the wave-lengths of the various lines in 
the spectrum are taken from the most accurate measurements: 


Ree PUMPING § 5 sia bs os be Rew en Rbs =795.0 up 

SRP MMO Gr a bes eh eens seo Sw ae ales Rb, =781.1 up 
j potassium. double line..-....2......... K = 166.9 mp 
SL OO. Oe 
: SEMA T So re. eee a poe eal s Li, =670.8 up 
| RCA NO oe hi elon oad eae es Cd, =643.9 up 
Orange-yellow lithium line................ Lig +=610.3 pp 

Yellow sodium (middle of the double line)...Na  =589.3 up 

RaPOHEY CAICWINL, DIG. f. )s os 6k bate ade Ss as Cag =554.4 pu 

MRSMR TEED TING 9. ihe 5. 0 yes ex wsatig Sac ¢ bes 83020 un 

MAMIDC TO Ce eee a eg Cd, =508.6 pu 

Roribe COMM HOF os ce ek oie eco heen sha Cdg =480.0 uu 
DPIC Boe he obs Ob sd thle Mate Cd, =467.8 up 

SPOR GRNAIEL PIMOS ogg enka) db eh sone eras eds Sr; ~=460.7 wu 

RAPOUMISR UTNE igs So ain ele! Su Sap pane ied Csg =459.3 uy 

OE oo SUSE aS ee ial Ree Cs,  .=455:3. ap 

Blue-violet indium line...............:... Ing.) = 451; Tae 

WAGIEL PMDICUUIN TING. Go) eek es ek dood Rbg =421.5 up 

WER RCRRIS SURI te PGs oaks he 60s Rb, =420.2 up 

BEA TEI Bi GS cons eae Re headed sos Ing =410.1 pp 

WMIDGAMIIIN TO Ko Sol oes slere ceic es Kg, =404.4 up 

mp (i ranennorer Gln)... 6355200 asco wcns ss = 396.8 up 


Let us assume that the foregoing lines were observed at the following 
positions on the scale: K, at 17, Kg at 154, Li, at 32, Na at 50, Tl at 68, 
Sr; at 101, In, at 111, and Ingat 149. | 


* The cadmium lines can be seen distinctly only in spark spectra. 





118 


‘ 


REACTIONS OF THE METALS 


If now, upon a rectangular system of coordinates, we plot the wave 
lengths as ordinates and the corresponding scale divisions as abscissas, 











and connect the points of intersection, 


























ad ae we shall get a curve* which expresses 
"75" os the relation between the wave lengths 

and the corresponding divisions on the 
we scale (Fig. 9). We can, without serious 
125-4 error, regard the portions of the curve 

between two points of intersection as 
wach proportional to the corresponding scale - 
675- divisions. 

Lia The use of the curve may be illus- 
rd trated by a single example: 
625- Suppose we observe at the division 

60 of the scale a linex. What is the 
gi corresponding wave length? The line 
BT5- ge lies between the sodium and thallium 
oe lines, which are of known wave lengths. 

Tla 
525- 
500- 
Hp 

fi 
450 Ing 
ag me Kp 
400 } | tT t I T T T T - — 

20 40 50 60 70 80 90 100 110 120 130 140 150 160 














Fia. 9, 


The sodium line lies at division 50 and corresponds to 589.3 pu 
The thallium line 


68 535.0 pu 


Consequently 18 scale divisions= 54.3 up 


and lscale division=3.017 py 


The line x is at 60=50+10 divisions, and these correspond to 
Division 50=589.30 yu 
10 divisions= 30.17 py 


Division 60= 559.13 wu 


* The curve is ordinarily drawn only for wave lengths lying between 800-400 pp 
because only these light rays are visible to the eye. The rays of shorter wave 
length than 400 pp are called ultra-violet rays and those longer than 800 py are 
The former can be detected by their action upon photo- 


known as ultra-red rays. 


graphic plates, and the latter by their thermic effect or by their action upon 
specially prepared photographic plates. 





; 
kL. 


BARIUM 119 
The scale divisions increase in their value as the wave lengths 
diminish, so that we subtract the 30.17 from 589.3. 


The spectra thus obtained are very beautiful, but usually of short duration. 
More permanent spectra may be obtained with the aid of the Beckmann burner, * 
Fig. 8, p. 116. From one-half to 1 cc. of the salt solution to be tested is placed in 
the glass vessel G, the gas is lighted, and a good blast of air introduced at a. By 
means of the air current, a little of the solution is carried mechanically over into 
the burner in the form of spray, and thus the salt reaches the non-luminous flame. 
In this way it is possible to get a fairly permanent spectrum by the use of only a 
few milligrams of substance. There is then plenty of time to make the measure- 
ments without having to stop and replenish the sample, 


Measuring the Lines and Bands of the Spectrum 


The correctness of a curve of wave-lengths, prepared as just 
described, depends upon the accuracy with which the observed lines 
or bands are measured. ‘To insure accuracy, all the better forms of 
apparatus are provided with cross-hairs in the ocular, and the cross must 
each time be made to cover a definite part of the line to be measured. 
The choice of such position depends upon the construction of the slit 
at the end of the spectroscope. If the apparatus is provided with an 
unsymmetrical slit, i.e., one of which only one side is movable while the 
other remains in a fixed position, then the reading should be made with 
the cross-hairs coinciding with the edge of the band on the immovable 
side of the slit. Such a position is shown in Fig. 10, in which it is 
assumed that the right-hand side of the slit is immovable. 

In the case of a symmetrical slit, i.e., one in which both sides of 
the slit open and close symmetrically, the cross-hairs must be made to 
meet in the center of the line to be measured (Fig. 11). 

The measurements can be made more readily and more accurately 
by the aid of Hilger’s Wave Length Spectroscope (Fig. 12). In this 
apparatus the telescope and collimator are both fixed in position, but 
the prism can be rotated by means of a cylinder upon which is engraved 
a very exact wave-length curve (Fig. 13). The pointer of the cylinder 
gives the desired wave length with an accuracy of +0.2 uy. 

For the rapid adjustment of the apparatus, the ocular shown in 
Fig. 14 is convenient. In this ocular there is a very fine, polished steel 
point beneath the cross-hairs; it is lighted from the outside by means of 
a small mirror so that it is seen very distinctly. This point is made to 
coincide with the desired edge, or the middle of the line, and then only 
the final adjustment is made with the cross-hairs. 





* Cf. Z. phys. Chem., 11, 472. 


120 ) ‘REACTIONS OF THE METALS 


It is not possible to volatilize all substances in the gas flame. The. 
non-volatile substances will not give any flame spectra, but by means of 
the electric spark they will be volatilized enough to give spark spectra. 
As the determination of their spark spectra furnishes the only method 
by which the purity of many substances may be tested, the methods for — 
the production of spark spectra will be outlined briefly. 

The apparatus invented by Bunsen consists of two platinum wires 
which are each attached at one end to conical pieces of charcoal. The 
latter are soaked with a solution of the substance to be tested. The 





Movable Movable 


Y YW 
J SAIN 


es 

















Fic. 10.—Position with un- Fig. 11.—Position with sym- 
symmetrical slit. metrical slit. 


two carbon points are now placed opposite to one another, quite near 
together, and the other ends of the platinum wires are connected with 
an induction coil, which causes sparks to pass between the two carbon 
points, volatilizing some of the salt. If now the sparks are viewed 
through a spectroscope, a large number of lines will be seen, of which 
only a part are produced by the substance itself. Some of the lines are’ 
caused by the carbon points and some by the air. In order to determine 
which lines are caused by the substance that is being tested, the experi- 
ment is first performed without using any of the substance, and the 
spectrum thus obtained is either drawn or photographed. The experi- 
ment is then performed with some of the salt impregnated in the carbon 
points and the new lines in the spectrum will be caused by the substance. 
Spark spectra can be more conveniently produced by means of the 





BARIUM 121 


so-called “ fulgurator ”’ of Delachanel and Mermet, as shown in Fig. 15. 
‘The salt solution is in a test-tube, so that the slit of the spectroscope 


cannot be contaminated with the spattered particles of the salt. 







Massegeatl! 
j 
| 


il 









Lh 
iS, 


U 





| LAAT 
LT 























ro) 
fl 63 
eee, 
3 ee 
= 
385 


i CT 


FOTO TLL 
il 





Fria. 14.—Ocular with steel point 


Fic. 13.—Cylinder (enlarged). 
and mirror. 


The small apparatus of H. Dennis, Fig. 16, is also very convenient 
for producing spark spectra. The platinum wire z is fused in a glass 
tube and ends in a point of Ceylon graphite, which extends up out of the 
arm e of the apparatus. The upper pole is not shown in the illustration. 


122 REACTIONS OF THE METALS 


To fill the tube m with the salt solution of the substance to be tested, 
sis taken away, and the apparatus is inclined to the left. The solution 
is poured in at the upper end of m, when s is again introduced and shoved 
down until it reaches nearly to the bottom. The apparatus is now 
placed in a vertical position, whereby the liquid in e rises until it is 
level with the lower end of s. On raising s the liquid will rise to the 
upper edge of e. By means of electric sparks from the carbon point, 
enough of the liquid is evaporated to give the desired spectrum. Care 
should be taken to prevent any spattering 
of the substance into the slit of the spec- 
troscope. 

A far better fulgurator has been devised 
by E. H. Riesenfeld and G. Pfiitzer.* As 
anode the solution of the salt itself is used, 
as cathode a thin iridium wire which almost 
touches the surface of the liquid. If an 
electric are is produced between such elec- 
oes trodes, nearly pure spectra are obtained 
: without lines from the atmosphere. In this 
way even the magnesium lines can be 
identified. By flame spectra with the 
Bunsen burner only 0.2 mg. Ca, 0.6 mg. Sr 
and 14mgs. Ba can be detected with cer- 
tainty, but with this fulgurator 0.002 mg. 
Ca, 0.03 mg. Sr and 0.006 mg. Ba can be 
detected when present in 1 cc. of water. 

To examine gas spectra, small Geissler 
tubes are used which contain the gas to be 

Fra. 15. Fia. 16. detected in a dilute condition. Besides 

flame spectra and spark spectra, there re- 
main absorption spectra to be mentioned. If white light is passed 
through a colored solution or gas, certain rays are absorbed by the 
liquid, so that if the light is now examined in the spectroscope, — 
these rays will be found to be lacking. We see a bright spectrum 
broken by black bands (absorption bands) which are character- 
istic for different substances. Thus solutions of permanganic acid, 
neodymium, praseodymium, erbium, and many other substances give 
characteristic absorption spectra. These absorption spectra are often 
obtained from colorless solutions. 

It is to be noted that an absorption gaseunten is often of quite 
different appearance according to whether the solution is dilute or 


* Ber., 1918, 3140. 





| 






































BARIUM 123 


concentrated. In measuring absorption bands the solution should 
be diluted with solvent until the band appears in the form of as fine a 
line as possible at the intersection of the cross-hairs, and further dilu- 
tion should cause the disappearance of the band. 

The same end is easily attained if, as R. Philip has suggested, the 
concentrated solution is placed in a test-tube and covered with more 


: a aa BERSE 
PSSST SSS 














s TT HHH 
N SS 

BSS SSS SS y 
ESS SSSR MSS 


Fic. 17. 





of the solvent. The lowest layer then gives the spectrum of the con- 
centrated solution and it becomes more and more dilute in the upper 
portions. 

_ In order to obtain any of the above spectra sharply defined, it is 
necessary to have the spectroscope properly adjusted, i.e., the rays must 
come parallel from the collimator tube into the telescope. The telescope 





is removed and adjusted for parallel rays by focusing it upon some 
distant object. It is then replaced, the prism removed, and the colli- 
mator tube brought exactly opposite the telescope, so that the slit of 
the former can be observed by the latter. The collimator tube must be 
- lengthened or shortened until the picture of the slit is sharply defined. 
The prism is now replaced and the scale tube adjusted until the scale 
also can be seen clearly defined. The instrument is now ready for use. 

The direct-vision pocket spectroscope, with an arrangement of 
prisms as shown in Fig. 17, is very convenient for ordinary use 


124 REACTIONS OF THE METALS 


If it is desired to photograph a spectrum, the telescope tube is 
replaced by a camera as shown in Fig. 18. Such an apparatus is 
called the spectrograph. With glass prism and glass objective lens, 
the spectrum can be photographed to about 350 yu. To make the 
ultraviolet part of the spectrum visible, the objective lens and the prism 
must be prepared of quartz and fluorspar, whereby it is possible to 
photograph light rays of about 200 yu. For observing rays of still 
shorter wave length, fluorspar alone is requisite, as it absorbs the ultra 
violet rays to a less extent than quartz does. 

To photograph the ultra red rays, specially prepared plates, as dis- 
covered and prepared by W. Abney, are necessary.* 





*For methods of quantitative spectrum analysis consult W. Hempel and R. L. 
von Klemperer, Z. angew. Chem., 1910, 1756. 


“aa 


GROUP III. AMMONIUM SULFIDE GROUP 


ALUMINIUM, TITANIUM, CHROMIUM, IRON, URANIUM, ZING, 
MANGANESE, NICKEL, COBALT 


ALUMINIUM, Al. At. Wt. 27.1 
Sp. Gr. 2.56-2.67. M. Pt. =658.7° C. 


Occurrence.—Aluminium occurs very extensively in nature, prin- 
cipally in the form of silicates, of which the feldspars and micas with 
their decomposition products are important examples: 

Orthoclase (feldspar), KsAls(Sig0g)3; muscovite (mica), 
KHoeAl3(Si04)3; kaolin (decomposed feldspar), H4aAleSizO9. Impure 
kaolin is called clay. 

Among the most important minerals which contain aluminium 
may be mentioned cryolite, NagAlF¢; spinel, MgAleO4, or magnesium 
aluminate, which crystallizes in the regular system and is isomorph- 
ous with magnetite, Fe-Fe2O4, and chromite, Fe-Cr204; alunite, 
KAI3(SO4)2(OH)6; aluminium hydroxide as hydrargillite, Al(OH)s, 
monoclinic; bauxite, H4AleO5, and diaspore, HAIOg, orthorhombic. 

Corundum, AlgOs, is next to the diamond the hardest mineral. 
Pure, transparent corundum often has a. beautiful color and is classed 
with the precious stones. Thus the ruby, sapphire, oriental emerald, 
oriental topaz and amethyst are nearly pure corundum colored by a 
little foreign oxide. Emery is an intimate mixture of corundum, 
magnetite and iron sulfide. True topaz is Ale(F,OH)2S8i04, turquoise 
is Alo(OH)3PO4-H20, and garnet is chiefly CagAle(Si04)s. . 

Metallic aluminium is now made on a large scale by the electrolysis 
of AlsOs dissolved in a bath of molten cryolite. It has a silver-white 
color and is only slightly attacked by exposure to the atmosphere. 
From the position of this element in the electromotive series, (p. 41) 
the metal should be attacked readily by the atmosphere, and the reason 
it is apparently unattacked has been proved due to its becoming 
coated with a thin, firmly adherent, protective layer of oxide. Alumin- 
ium is trivalent in all its compounds; it forms only one oxide, Al2Os, 
which is amphoteric. The metal dissolves in acid to form an aluminium 
salt and in caustic alkali to form a soluble aluminate. Aluminium 
readily replaces the hydrogen of hydrochloric acid, but it dissolves 

125 


126 REACTIONS OF THE METALS 


less readily in dilute sulfuric acid and becomes passive when treated 
with nitric acid. One theory of the cause of passivity is the formation 
of a closely adherent oxide film. 

The action of aluminium upon dilute hydrochloric acid and upon 
aqueous solutions of caustic alkali is expressed by the following ionic 
equations: 


2A1+6H* — 2Al***+3H:2 T, 
2A14+20H-+2H20 > 2Al02-+3He Tt. 


Aluminium salts are as a rule colorless, and those which are soluble 
in water show an acid reaction in aqueous solution, on account of their 
being hydrolyzed to a considerable extent. This explains the fact 
that on evaporating a solution of aluminium chloride in water we do 
not obtain aluminium chloride, but the insoluble oxide, or hydroxide: 


AICl3-+3HOH @ Al(OH)3+3HCI. 


The property which aluminium possesses of forming alums is very 
characteristic. The alums are double salts of aluminium sulfate with 
the sulfates of potassium, cesium, rubidium, or ammonium, of the 
general formula RAI(SO4)2:12H2O, in which R represents one of the 
univalent metals just mentioned. Similar salts in which the aluminium 
is replaced by trivalent iron or chromium are also called alums. The 
alums crystallize in the regular system, usually in octahedrons. The 
common potassium alum is much less soluble in cold than in hot water. 
Thus 100 cc. of water at 15° C. dissolve 10.7 gms. of alum and 283 gms. 
at 100° C. 

The sulfide of aluminium can be prepared only in the dry way. 
It is a pale-yellow substance, which is hydrolytically decomposed even 
by cold water into hydrogen sulfide and aluminium hydroxide; 


AloS3+6HOH =3HeS 7 +2Al1(OH)s. 
Toward strong: acids aluminium hydroxide plays the part of an 
alkali, while toward strong bases, on the other hand, it acts as an acid. 
REACTIONS IN THE WET WAY 


1. Ammonia produces a gelatinous precipitate of aluminium hydrox- 
ide, which is somewhat soluble in water, but insoluble in the presence of 
ammonium salts: 


AICl3+3NHs0OH = Al(OH)3+3NH4Cl. 





ALUMINIUM 127 


The property which the aluminium hydroxide shows of partly 
dissolving in water, is common to all colloidal substances (cf. p. 59). 
When dissolved they are sometimes said to exist in the hydrosole con- 
dition and when precipitated as hydrogele. : 

The hydrosole form of aluminium hydroxide can be converted into 
hydrogele by the addition of salts,* preferably ammonium salts. [Jf, 
therefore, we desire to precipitate aluminium from a solution by means of 
ammonia, we take care that ammonium chloride is present. } 

The freshly precipitated aluminium hydroxide is readily soluble in 
dilute acids; but after standing some time in a salt solution, or after 
long boiling, it becomes more difficultly soluble, so that it is necessary 
to digest it with acid for a long time in order to bring it completely into 
- solution. 

The solubility product of aluminium hydroxide is so small that it is 
precipitated by ammonia even in the presence of ammonium salt. 
Aluminium does not show a tendency to form soluble complex cations 
with ammonia. | 

2. Potassium or Sodium Hydroxide produces the same _ precip- 
itate as ammonia, which is, however, in this case soluble in excess of 
the reagent, forming an alkali aluminate: 


Alt++430H- > Al(OH)s, 
Al(OH)3+OH~ — AlO2~+2H20. 


If dilute acid is added to a solution of an aluminate, there is formed 
at first a precipitate of aluminium hydroxide, which dissolves on further 
addition of acid: 


AlO2- +H*+H20 — Al(OH)s, 


Al(OH)3+3H+ > Al**+*+3H20. 


The aluminates are also decomposed by boiling with an ammonium 
salt, 


AlOo-+NH,4t+H20 > Al(OH)3+NHs3 tT ‘ 


Aluminium hydroxide is soluble in neutral tartrates of the alka- 
lies, so that in the presence of tartaric acid there will be no precipi- 
tation on the addition of ammonia. Consequently the aluminium 





* This principle is illustrated by the technical process of “ salting out ’’ colloidal 
dyes from their solutions.: 


128 REACTIONS OF THE METALS 


cannot be present in the solution in the form of the simple alumin- 
ium cation, but as a complex negative ion: 


Al(OH)3+CsH406 — C4Heo(AlIOH)Og +2H20. 


Many other organic hydroxy-acids and other hydroxy-compounds, 
such as malic and citric acids, sugars and starches, have the same 
effect of preventing the precipitation of aluminium hydroxide by am- 
monia. 

3. Ammonium Sulfide causes precipitation of the hydroxide. 
When formed in the dry way, aluminium sulfide is comp'etely hydrolized 
by water: 

Al.S3-+6HOH =2Al(OH)3+3H2S 1. 


The aqueous solution of ammonium sulfide is in equilibrium with 
NH4OH and HeS formed by hydrolysis. There are, therefore, enough 
OH ions in a solution of ammonium sulfide to cause the precipitation 
of Al(OH)s. 

4. Alkali Carbonates precipitate aluminium hydroxide (hydrolysis): 


2AICls f 3NaeCO3 =6NaCl +Alo (COs3)s3, 
Aleo(CO3s)3-+6HOH = 3H2CO3+2Al(OH)s3. 
— 
3(H20+COsz T ) 


5. Barium Carbonate, suspended in water and added to the solu- 
tion of an aluminium salt, also precipitates the hydroxide: 


2AICl3 +3BaCO3+6HOH =3BaCle+3H2CO3+2Al(OH)3. 


—— . 


3(H20+CO0z2 1) 


6. Alkali Acetates produce no precipitation in cold neutral solu- 
tions, but, on boiling the solution, a very voluminous precipitate of 
basic aluminium acetate is formed: 


AICl3 +3NaC2H302 =3NaCl+ Al(C2H302)3 (in the cold), 


Soluble 
SO me 
_ Al(C2H302)3+2HOH @ mice +2HC2H30¢2 (on boiling). 
CoH302 


If the solution is allowed to cool, the basic aluminium acetate 
redissolves. The reaction is, therefore, a reversible one, and goes 
more completely from left to right according as we increase the amount 
of water and raise the temperature (cf. pp. 50 and 51). 





ALUMINIUM 129 


7. Alkali Phosphates, e.g., NasHPOu,, give a gelatinous precipitate 
of aluminium phosphate: 


2HPO4« +Al**+ — AIPO,+HePOz, 


or, on addition of ammonia, 


HPO.-+NH3+Al*+t+ — AIPO,+NH,". 


Aluminium phosphate is soluble in mineral acids, insoluble in 
acetic acid (differing from Ca, Sr, Ba, Mg), but readily soluble in 
sodium or potassium hydroxide solutions: 


AIPO.,+40H — AlOo + PO4"+2H20. 


On boiling this alkaline solution (obtained in the last reaction) 
with ammonium chloride, a precipitate will be formed, consisting 
of a mixture of aluminium phosphate and aluminium hydroxide; 
while barium chloride, on the contrary, will precipitate barium phos- 
phate from such a solution and leave the aluminate dissolved. 

8. Sodium Thiosulfate, NazS2O0s,. completely precipitates the 
aluminium as hydroxide on boiling: 


2Al***+38203"+3HOH =38S+3S802 T +2A'(OH)s. 


The sodium thiosulfate, being a salt of a weak acid, serves to neu- 
tralize the H* ions formed by the hydrolysis of the aluminium salt, and 
the free H2S203 is so unstable that it breaks down into S and SOz; 
equilibrium is disturbed by the precipitation of the Al(OH)3 and § 
and by the escape of SOz gas. 

9. Morin in alcoholic solution shows a green fluorescence when 
brought in contact with only a trace of neutral aluminium salt. (This 
reaction is very sensitive.) Beryllium salts and salts of the rare earths 
do not give the test. 

10. Ether precipitates white, crystalline aluminium chloride, 
AICl3-6H2O, from a concentrated solution which is saturated with 
HCl gas. Aluminium may be separated from beryllium in this way. 


DETECTION OF ALUMINIUM IN THE PRESENCE OF ORGANIC SUBSTANCES 
WHICH PREVENT THE PRECIPITATION BY THE ABOVE REAGENTS 


The presence of tartaric acid or other non-volatile, organic hydroxy-com- 
pound prevents the precipitation with above reagents. To detect the presence 
of aluminium in such cases, the organic substance must be first destroyed. This 
is best accomplished as follows: Add some sodium carbonate and a little 


130 REACTIONS OF THE METALS 


potassium nitrate to the solution, evaporate to dryness in a platinum dish, and 
ignite the residue, whereby the aluminium becomes aluminate and the organic 
substance is destroyed with separation of carbon. Treat the residue with dilute 
nitric acid and filter; the aluminium goes into solution as nitrate and can 
be precipitated with any of the above reagents. 

If sufficient nitrate was present, the carbon will be completely burnt to CO,, 
and the residue then often contains undecomposed nitrate or nitrite. It would, 
therefore, be unwise to treat the residue in the platinum dish with hydrochlorie 
acid, as aqua regia would be formed and the platinum would be dissolved. Con- 
sequently the residue is treated with nitric acid (or with hydrochloric acid in a 
porcelain dish). : 


When strongly ignited, aluminium hydroxide loses water and 
forms the anhydride, Al2O3, which is scarcely soluble at all in hydro- 
chloric and nitric acids. In hot, concentrated sulfuric acid, with 
a little water, it will dissolve after long digestion. The ignited alumin- 
ium oxide, as well as the natural corundum, is most readily brought 
into solution by fusion with potassium pyrosulfate. The fusion is 
accomplished in the following way: Take twelve times as much fused 
potassium acid sulfate as there is oxide to get into solution and heat it 
by itself in a large platinum crucible over a small flame. The acid 
potassium sulfate melts readily, at about 300° C., gives off water 
(causing frothing), and becomes changed into potassium pyrosulfate: 


2KHSO4 => HeO0O+ KoSe07. 


As soon as the frothing has ceased, the transformation into potas- 
sium pyrosulfate is complete. Add the dry aluminium oxide to the 
crucible and continue heating until the melt begins to solidify (showing 
that a considerable amount of potassium sulfate, which melts much 
more difficultly than the pyrosulfate, has been formed), then raise the 
temperature and continue heating until the oxide has dissolved clear in 
the melt. By heating the pyrosulfate, SO3 is given off, which at the 
high temperature is very active: 


K28207 = Ke804+80s3. 


If the heating is too rapid, much of the SOsz is lost. 
After the reaction is finished, the melt contains the aluminium 
as aluminium sulfate in the presence of potassium sulfate, 


3Ko8207+ AloO3 = Alo(SO4)3-+3K280a, 


and both of these substances can be brought into solution by treating 
them with water. 


ALUMINIUM 131 


The ignited aluminium oxide can also be brought into solution 
by fusion with caustic alkalies: 


AlzO3+2KOH = 2KA102+H20. 


This last operation is usually carried on in a silver crucible, never in 
platinum, as the latter would be strongly attacked. 

Native AlzO3 (corundum, ruby, sapphire, emery) can be completely 
brought into solution by fusion with potassium pyrosulfate. 


REACTIONS IN THE DRY WAY 


Aluminium compounds, on being heated with sodium carbonate 
on charcoal before the blowpipe give a white, infusible, brightly glowing 
oxide, which, when moistened with cobalt nitrate solution and again 
heated, becomes a blue infusible mass (Thénard’s blue) : 


AleO3 +CoO = CoAleQOx.. 


In carrying out this test it is extremely important not to use an 
excess of cobalt nitrate, for this salt leaves black cobalt oxide behind:on 
ignition and when an excess is present it entirely obscures the blue color 
of cobaltous aluminate. 

The test is usually applied to a precipitate of aluminium hydroxide. 
The best way to carry out the test is as follows: Dissolve the precipitate 
in 5 ec. of 6 N nitric acid and add to the solution half as many small 
drops of 1 per cent cobalt nitrate solution as there are presumable 
milligrams of aluminium present. Evaporate the solution nearly to 
dryness, add a few drops of water and moisten a roll of filter paper 
with the concentrated solution. Heat over the flame until the paper 
is entirely consumed and then very strongly. 

Thénard’s blue is infusible. If a fused mass is obtained in this test, 
the presence of aluminium is doubtful, because all fused glasses, such as 
borax beads, sodium phosphate beads, etc., are colored blue by cobalt 
oxide. 

Aluminium salts are not volatile, and do not color the flame. By 
ignition in the air, all aluminium salts, with the exception of the phos- 
phate and silicate, are decomposed, leaving behind the oxide: 


PAICI3 +30 = AloO3+3Cle T , 
2Al(NO3)3=Als03+3N20s 1 , 
Alo(SO4)3 = Als03+3803 T . 


132 REACTIONS OF THE METALS 


CHROMIUM, Cr. At. Wt. 52.0 
Sp. Gr.=6.81. M. Pt.=1510° C. 


Occurrence-—Chromium occurs in nature as chromite, FeCreOs, 
isomorphous with spinel (see aluminium); as monoclinic crocoite, 
PbCrO4; and as laxmannite, a double compound of lead-copper phos- 
phate and basic lead chromate, (Pb,Cu)3(PO4)2-PbsO0(CrO4)o. Fur- 
thermore, it is found in small quantities in many silicates, such as 
muscovites, biotites, augites, etc., and consequently in their weathering 
products, as in many kaolins, bauxite, etc. 

Metallic chromium is a white, crystalline powder. With oxygen 
it forms the following oxides: chromous oxide, CrO; chromic oxide, 
Cr2O3; chromic acid anhydride, CrO3; and chromium peroxides 
corresponding to CrO4, CreOg9, CreO11 and CreOj3 (ef. p. 139). 

The oxides CrO and Cr2O3 are basic anhydrides, and, on being 
dissolved in acids, yield the corresponding salts, the chromous and 
the chromic compounds. Chromium trioxide is the anhydride of the 
hypothetical chromic acid, HeCrO4 and forms chromates with bases. 
The chromium peroxides have never been obtained pure; salts of per- 
chromic acid have been isolated and analyzed. (Cf, p, 139.) 


I. Chromous Compounds 


Chromous oxide is known only in the form of its hydroxide, Cr(OH)s, 
which, on being dried, loses hydrogen and water, leaving behind chromic 
oxide: 


2Cr(OH)2 = He +H2O + CreOz. 


Like chromous hydroxide, all chromous compounds are extremely 
unstable, being changed readily by contact with the air into chromic 
compounds. Only the halogen compounds, the phosphate, carbonate, and 
acetate, are known in the dry state; the sulfate only in solution. ‘The 
acetate, Cr(C2H302)2+H20, is a reddish-brown, crystalline substance, 
insoluble in water, but readily soluble in hydrochloric acid. This 
solution, as well as that of all chromous salts, absorbs oxygen with 
avidity, and is consequently used in gas analysis for the determination 
of oxygen in gas mixtures. Solutions of chromous compounds) are 
obtained by the reduction of chromic compounds with nascent hydrogen 
(zine and acid), out of contact with the air. 

On account of the instability of these compounds the ariabplical 
chemist will rarely meet them and further description is unnecessary. 





: 


CHROMIUM 133 


II. Chromic Compounds 


All chromic compounds contain chromium as a trivalent element; 
they are colored either green or violet, and are soluble in water as 
a rule. The oxide, hydroxide, phosphate, anhydrous chloride, and 
the sulfate, after being strongly heated in a stream of carbon dioxide 


gas, are insoluble in water. Violet chromium chloride obtained in 


the dry way is insoluble in acids; it dissolves readily in water containing 
a trace of chromous chloride, or in the presence of stannous chloride 
(tin and hydrochloric acid). By dissolving the grayish-green chromic 
hydroxide in acids, green solutions are always obtained, which on 
long standing become greenish violet or violet, but on boiling become 
green again. Chromic sulfate forms with sulfates of potassium, 
ammonium, cesium, or rubidium, the so-called chrome-alums, which 
crystallize in the regular system. ‘These alums, like all other chromic 
salts, react acid in aqueous solution (hydrolysis). 

Chromic sulfide, CreS3, can be obtained only in the dry way. 
On being treated with water it is decomposed quantitatively into 
hydroxide and hydrogen sulfide. 


REACTIONS IN THE WET WAY 


1. Ammonia produces a erayish-green, gelatinous precipitate of 
chromic hydroxide: 
Crt +++30H- — Cr(OH)s. 


Chromic hydroxide is somewhat soluble in excess of ammonia, form- 
ing a violet or pink solution, particularly soluble when the ammonia 
is added to a violet solution of a chromic salt, in the presence of am- 
monium salts. This is caused by the formation of complex chromic- 
ammonia cations, which, however, may be decomposed by boiling 
the solution until the excess of ammonia has been driven off, when the 
chromium is quantitatively precipitated as hydroxide. In order, then, 
to precipitate the chromium from a solution as hydroxide, it is neces- 
sary to precipitate at a boiling temperature, and to use as little 
ammonia as possible. | 


By ignition of chromic hydroxide, green chromic oxide is obtained, which 
after strong ignition is insoluble in acids. In order to bring it into solution, it 
must be fused with potassium pyrosulfate (cf. aluminium, p. 130); or with 
an oxidizing flux such as sodium carbonate and potassium nitrate in a platinum 
crucible or sodium peroxide in a nickel crucible, whereby sodium chromate is 
formed: 

| Cr.0;+2Na.CO;+30 =2Na,Cr0O,+2CO, T “3 


134 REACTIONS OF THE METALS 


If the product of this last fusion is dissolved in water, acidified with hydro- 
chloric acid, and boiled with alcohol, a green solution of chromic chloride will 
be obtained (p. 188), from which the chromium can be precipitated as 
hydroxide with ammonia. On fusing with sodium carbonate and potassium 
nitrate in a platinum crucible, the crucible will always be slightly attacked, so 
that a small amount of platinum will go into solution with the fused mass; it 
can be removed, after the treatment with hydrochloric acid, by passing hydrogen 
sulfide into the boiling solution, and filtermg off the precipitated platinum 
sulfide. 


2. Potassium and Sodium Hydroxides cause the same precipitation 
as ammonia; but the precipitate is readily soluble in excess of the 
reagent, forming a green chromite: 


Crt +t+4+30H— —> Cr(OH)s3; Cr(OH)3+OH 2 CrO.-+2H20. 


Chromic hydroxide behaves here as a weak acid. The. reaction 
is reversible, the presence of considerable water causing the reaction 
to go from right to left, particularly at the boiling temperature. By 
boiling the dilute solution, complete hydrolysis takes place; the chro- 
mium is almost quantitatively precipitated as hydroxide (differing from 
aluminium). 

Chromic hydroxide often causes the precipitation of other bases 
as insoluble chromites, particularly zinc and magnesium. 

3. Sodium Peroxide and caustic alkali added to the solution of a 
chromic salt converts the trivalent chromium into the chromate ion | 
in which the chromium has a positive valence of six: 


2Cr* ++43Na002+40H- — 2CrO4-+6Nat+2H.0. 


4, Alkali Carbonates, Barium Carbonate, Ammonium Sulfide, 
and Alkali Thiosulfates precipitate chromic hydroxide, as with 
aluminium. 

5. Alkali Phosphates give a greenish, amorphous precipitate of 
chromic phosphate: 


2QHPO4-+ Crt t+* — CrPO4+HePO,". 


Chromic phosphate is readily soluble in mineral acids and in cold 
acetic acid. On boiling.the acetic acid solution, chromic phosphate 
separates out again. 

6. Alkali Acetates produce in solutions of chromic salts no pre- 
cipitation, even when the solutions are boiled. If, however, consider- 
able amounts of aluminium and ferric salts are present at the same 
time, the chromium will be precipitated almost quantitatively with the 
iron and aluminium as basic acetate. In case, however, chromium 
predominates, only a part of the metals will be precipitated as basic 





CHROMIUM | 135 


salts; the filtrate will contain iron and aluminium with chromium. 
In the presence of. chromium, the basic acetate separation is always uncer- 
tain. 

7. Hydrogen Sulfide produces no precipitate in acid solutions 
of chromic salts. 


III. Chromates 


Chromium trioxide, CrOzs, forms red orthorhombic needles, which 
dissolve readily in water to an orange-red solution. If this solution 


is neutralized with potassium hydroxide, it becomes yellow, and on 


evaporation yellow KeCrO,, the potassium salt of chromic acid, H2CrOu, 
is obtained. If the yellow solution of potassium chromate is acidified, 
and then allowed to crystallize, orange-red prisms of triclinic potassium 
dichromate crystals, Ko2Cr207, are formed. 

The aqueous solution of potassium chromate, K2CrO4, contains 
colorless potassium ions and yellow CrO4~ anions, while the aqueous 
solution of potassium dichromate, KeCreO7, contains, in the presence 
of the colorless potassium ions, the orange-red colored Cr207~ anions: 


KoCre07 = 2Kt + CreO7- - 


We are able, therefore, to determine from the color of a chromate 
solution the nature of the chromate ion which is present. 


The free acids, H»CrO, and H.Cr,O;, cannot be isolated, but only the corre- 
sponding anhydride, CrO;, which is very soluble in water. When chromium 
trioxide dissolves in water, the following reaction takes place: 


CrO; +4- H.O @ H.CrO,. 


The first hydrogen of this acid belongs to the class of very strong acids (cf 
p. 10) and consequently the greater part of the H.CrO, undergoes ionization 
as fast asitisformed: H,CrO,— H*+HCrO,. Thesecond hydrogen, on the 


other hand, corresponds to that of a very weak acid and we may say that in 


the presence of hydrogen ions HCrO,~ is scarcely ionized at all. The HCr0,-, 
however, enters into the following equilibrium: 2HCr0O,- 2 H.O+Cr O;. 
The presence of hydrogen ions will evidently favor the formation of the dichro- 
mate ion, while dilution will favor the reverse reaction, and the presence of 
OH ions will cause the dichromate ions to disappear and yellow CrQ,- ions 
will take their place. 

Remark.—Although we may judge as to the color of the ions from the color 
of the solution, and often predict what the color of the solid salt will be, yet, on 
the other hand, we cannot tell what the color of the solution will be from that 
of the salt itself. Yellow lead iodide dissolves in water to a colorless solution, 
and the yellow and red iodides of mercury, although only silghtly soluble, also 
do not yield colored solutions. 

If the solution of a salt is colored, the salt itself will be colored; but the reverse 
ts not always true. 


/ 


136 _ REACTIONS OF THE METALS 


All chromates are insoluble in water, except those of the alkalies, 
calcium, strontium, and magnesium. All chromates dissolve in nitric 
acid, except fused lead chromate, which dissolves with difficulty. 


Formation of Chromates ‘ 


All chromium compounds may be readily oxidized to chromates. 
According to whether the compound is soluble in water or not, different 
methods are used to effect the oxidation. 


The student should not attempt to memorize a large number of chemical 
equations, but he should strive to become able to express his chemical knowledge 
in the form of equations or’ equilibria expressions. To balance equations 
representing the oxidation of chromium, it should be remembered that the 
chromium is changed from a positive valence of three (in the form of chromic 
cations or chromite anions) to a positive valence of six (in the form of chromic 
acid, chromate or dichromate ions). When the oxidation takes place by means 
of halogen, a halide is formed and the halogen is changed from the neutral 
condition to the form of a negative ion with unit charge. When hypochlorite 
is the oxidizing agent, the unit positive charge on the chlorine atom is lost and 
a unit negative charge takes its place, which corresponds to the loss of two 
positive charges. Similarly, lead peroxide and hydrogen peroxide have an oxidiz- 
ing power corresponding to the loss of two unit charges of positive electricity. 

The valence of the chromium in the anions CrO,-, CrO;- and Cr,O,- is 
found, in accordance with the rule given on p. 25, by subtracting the charge 
of the ion from the product obtained by multiplying the number of oxygen 
atoms present in the ion by its valence of two. In equilibrium expressions it 
is important to make sure that the algebraic sum of the positive and negative 
charges on one side is exactly the same as the algebraic sum of the charges on 
the other. With a little practice it is very easy to determine whether hydro- 
gen ions, hydroxyl ions or water molecules are required to make the equation 
balance. 


The oxidation in alkaline solutions is effected: 

(a) By the halogens. If sodium or potassium hydroxide is added 
in excess to a solution of a chromic salt, and chlorine or bromine is 
conducted into the solution, the oxidation will be complete in a few 
minutes: the green chromite becomes yellow chromate: 


2Cr02~+80H-+3Cle > 2CrO4"+6CI--+4H20. 


Chromic compounds may be also oxidized by halogens in the 
presence of sodium acetate, the reaction going extremely slowly in the 
cold, but very quickly on warming: 


2Cr* *++4+3Cle+8H20 & 2CrO4-+6Cl-+ 16H. 


The sodium acetate greatly lowers the concentration of the hydrogen 
ions and permits the reaction to proceed slowly in the direction left to 
right (cf. p. 46). 





CHROMIUM 137 


(b) By hypochlorites (sodium hypochlorite, chloride of lime, ete.): 
2CrO2-+30Cl +20H- — 2Cr04-+3Cl-+H20. 


(c) By lead peroxide. The alkaline solution is boiled with lead 
peroxide: 


2CrOe +3PbO02+80H — 3PbO2 +2Cr04-+4H20. 
(d) By hydrogen peroxide, 
2CrO2 +3H202 +20H” — 2CrO4 +4H20, 


the reaction taking place on warming. 
(e) By freshly precipitated manganese dioxide. The oxidation 
_ takes place on boiling the neutral or slightly acid solution: 


2Cr* +*+4+3Mn02+40H- — 2Cr047+3Mn* t+2H20. 


It is evident from this last equilibrium expression that the presence of. 


hydroxyl ions should favor the oxidation and hydrogen ions should hinder it. 
On the other hand, manganese dioxide is very insoluble in alkaline solutions. 
It is necessary, therefore, to have the solution neutral or slightly acid in order to 
obtain a sufficient concentration of quadrivalent manganese in solution. 

In acid solutions the chromic cation is the most stable condition for chrom- 
ium, but in alkaline solutions the chromate anion is the more stable condition. 
In acid solutions, therefore, it is easy to reduce a chromate to chromic salt and 
in alkaline solutions it is easy to oxidize a chromic salt to chromate. 


Oxidation in acid solution may be effected by boiling with very 
energetic oxidizing agents such as concentrated nitric acid and potassium 
chlorate, sodium bismuthate (or bismuth tetroxide), or potassium per- 
manganate: 


2Cr* +++3NaBi0g+4H* —> Cre07-+3Nat+3Bit +++2H20. 


In carrying out this reaction the chloride should not be used, or any other 
salt of which the anion is capable of oxidation; the acid used should 
be nitric or sulfuric acid. 

In the case of an insoluble chromium compound, such as strongly 
ignited chromic oxide, or the mineral chromite, the oxidation is effected 
by means of fusion with sodium carbonate and an oxidizing agent such 
as potassium nitrate, potassium chlorate or sodium peroxide (cf. pp. 
133, 142). The alkali chromates thus obtained are of a deep-yellow 
color, and are readily soluble in water. 


a Reduction of Chromates 


Chromic acid, chromates and dichromates are strong oxidizing 
agents in acid solutions. Such reactions often take place even in very 


~ 


/ 


138 REACTIONS OF THE METALS 


dilute solution, and for this reason potassium dichromate is often used 
in quantitative analysis, the quantity of reducing agent being de- 
termined by the volume of potassium dichromate required to react 
- with it. Ferrous ions, sulfurous acid, hydrogen sulfide and hydriodie 
acid are oxidized at the ordinary temperature. Oxalic acid and alcohol 
are oxidized slowly at the laboratory temperature and very quickly 
on heating the solution; hydrochloric acid and hydrobromic acid only 
when the solution is hot. The original orange solution is changed to 
green, the color of chromic ions; 


Cr207-+6Fe* ++14H* > 2Cr*++*++6Fet*+++7H20, 
Cr207- +3803" +8H*t — 2Cr++*+3S804-+4H20, 
Cr207-+3H2S+8H*t — 2Crt +*+3S8+7H20, 
Cr207~+3H2C201+8H* => 2Cr+*+*++3C02 T +7H20, 
Cr2O7-+6I- +14H* > 2Cr*+++3I1,+7H:20, 
Cr207~+6HC1+8H*t — 2Cr+*++3Cle T +7H20. 


As this last reaction takes place only on warming, it furnishes us 
with a convenient method for preparing small quantities of chlorine 
for analytical purposes, because the evolution of chlorine ceases as — 
soon as the lamp is taken away. It is necessary, however, to em- 
ploy an excess of hydrochloric acid, as otherwise no chlorine will be 
evolved owing to the formation of potassium chlorochromate, KCrOs3Cl: 


Cre07~+2HCl — 2CrOsCl + H20, 
which is decomposed on adding more hydrochloric acid: 


2CrO3Cl-+4HCI+8H*t > 2Crt+*+*++6H20+3Cl. T. 


If alcohol and hydrochloric acid are allowed to act simultaneously 
upon a chromate (the reaction takes place on gentle warming without 
the evolution of chlorine), the alcohol is oxidized to aldehyde: 

Cr207"+3C2Hs0H+8H* — 2Cr***+-7H20+ 3CHsCHO hy 
Idehyde 
This last reaction is often used for reducing a chromate, because the 
aldehyde (recognizable by its peculiar empyreumatic odor) and the 
excess of alcohol are easily removed by boiling the solution, and the 
latter then contains simply the chromium and the metal of the chromate 
as chlorides. 





a ec « 
: Se te 
J = | 


ee ee a a a 


CHROMIUM 139 


By boiling chromates with concentrated sulfuric acid, reduction 
takes place with evolution of oxygen: 


2KeCr207+8H 2804 = 2K2804+ 2Cre(SO4)3 +8H20+302 T. 


The behavior of free chromic acid toward hydrogen peroxide is 
characteristic. The chromic acid is converted into blue perchromic 
acids which are soluble in ether: H7CrOi0, H3CrO7 or HgCrOsg. 

If a cold, alkaline solution of a chromate is treated with neutral 
hydrogen peroxide, the solution is colored red, owing to the formation 
of an alkali salt of perchromic acid, H3CrOs: 


2KeCrO4+7H202+2KOH =8H20+2K3CrOsg. 


Little by little the red color disappears, with evolution of oxygen, 
and the yellow color of the chromate returns: 


4K3Cr0g+2H20 =4KOH+4KoCr04+702 fT. 


Tf a cold, neutral solution of potassium dichromate is treated with 
hydrogen peroxide, the solution is colored violet, due to the formation 
of the potassium salt of a slightly different perchromic acid, H3CrO7: 


KeCre07 + 5H2O0e2 = 3HeO -- 2K HeCrO7z. 


In this case, also, the violet color gradually disappears with evolution 
of oxygen and regeneration of the dichromate: 


4K He2CrO7 = 2KeCreO7 +502 T a 4H>0. 


If either the red or violet solution, obtained as above described, 
is shaken with ether, the latter remains colorless. 

The behavior of chromate solutions toward an excess of hydrogen 
peroxide in the presence of dilute sulfuric acid is quite different. 
There is then formed invariably the perchromic acid richest in oxygen, 
H7CrQjo, and the solution is turned an intense blue. The blue color dis- 
appears after a short time and the solution turns green, owing to the 
conversion of all the chromium into the chromic condition: 


2H7Cr019+3H2804 = Cro(SO4)3-+10H20+502 T . 


The perchromic acid is very soluble in ether; if, therefore, the 
aqueous solution is shaken with ether the latter becomes colored a 
beautiful blue. The perchromic acid is more stable in ethereal than 
in aqueous solution. 

Since the formation of the intense-blue-colored perchromic acid 
takes place so readily, it may be used as a basis for a sensitive test 


140 REACTIONS OF THE METALS 


for free chromic acid, which is made as follows: Add a few drops of 
dilute sulfuric acid to one or two cubic centimeters of hydrogen per- 
oxide and shake with 2 cc. of ether; then add a little of the chromate 
solution and shake the mixture again. In the presence of 0.1 mg. of 
chromic acid, the upper ether layer is colored intensely blue, and ‘the 
reaction is noticeable with only 0.007 mg. of chromic acid (Lehner). 

Most chromates are insoluble in water, and exhibit character- 
istic colors; therefore it is easiest to test for chromium when it is present 
as a chromate. 


Reactions for the Precipitation of Chromic Acid 


1. Sulfuric Acid.—Dilute sulfuric acid causes, at the most, a 
change of color from yellow to orange, without any evolution of gas. 

Concentrated sulfuric acid causes the cold solution to change 
to orange color, and there is often a separation of red needles of CrO3; 
the solution on being heated becomes green, the chromic acid being 
reduced to chromic salt with evolution of oxygen: 


4Cr03+6H2804 =6H20+302 7 + 2Cro(SOx)s. 


. 2. Silver Nitrate produces in neutral chromate solutions a brownish- 
red precipitate of silver chromate: 


CrO4g” + 2Agt —s AgoCrO4 ) 


soluble in ammonia and mineral acids (hydrochloric acid changes 
it into insoluble silver chloride and chromic acid), insoluble in acetic 
acid. If to a moderately concentrated solution of potassium dichro- 
mate, silver nitrate be added, a reddish-brown precipitate of silver 
dichromate is formed: . 


CroO7_+ 2Agt — AgeCreOz, 


which, on being boiled with water, is changed into the less soluble 
normal silver chromate: 


2AgeCr207 +H:20 — 2AgoCrO, + HeCre2QO7. 


The presence of sodium acetate causes this change to take place in 
the cold (cf. pp. 46 and 135). 

3. Lead Acetate produces in solutions of normal chromates and 
dichromates a ye'low precipitate of lead chromate, which is soluble 
in nitric acid but insoluble in acetic acid: 


CrO4” of Pbt ty PbCr O4 





CHROMIUM 141 


and 


Cr207~+2Pb(C2H302)2+H20 — 2HCeH302+ 2C2H3027 + 2PbCrOg. 


‘If lead nitrate is used instead of lead acetate, the precipitation is not 
complete unless sodium acetate is added. 

4. Barium Chloride produces in solutions of normal chromates 
a yellow precipitate of barium chromate: 


CrO4-+Ba*+t — BaCrOu, 


soluble in mineral acids, insoluble in acetic acid. From solutions 
of dichromates the precipitation is complete only on addition of ‘an 
alkali acetate (cf. p. 108). 

5. Mercurous Nitrate produces in the cold a brown, amorphous 
precipitate of mercurous chromate: 


CrO4” + Hgo(NO3)2—-Hg2Cr04+2NO037, 


which on being boiled becomes fiery-red and crystalline. 


Behavior of Chromium Trioxide and Chromates on Ignition 


Chromium trioxide is decomposed on ignition into chromic oxide 
and oxygen, 4CrO3=2Cr203+302 T. The chromates of ammonium 
and mercury behave quite similarly. Thus normal ammonium chromate 
on ignition is changed to chromic oxide, ammonia, nitrogen, and water. 
The reduction of the chromate is favored by the reducing action of 
ammonia which is present in excess. 


2(NH4)2CrO4=2HN3 7 +Ne2 1 +5H20 Tf +Cr20s. 
Ammonium Dichromate evolves only water and nitrogen: 
(NH4)2Cr207=4H20 T Nz T +CreOs. 


This decomposition takes place violently with scintillation. The 
chromic oxide which remains behind is very voluminous and reminds 
one of tea-leaves; consequently it is sometimes called ‘‘ tea-leaved 


chromic oxide.”’ : 
Mercurous Chromate is decomposed on ignition into chromic oxide, 
mercury vapors, and oxygen: 


AH g2CrO4 =2Cr203-+8Hg 1 +5027. 


The Dichromates of the Alkalies are changed on ignition into normal 
chromates, chromic oxide, and oxygen: 


4KoCre07 =4KoCr04+ 2Cre03+302 t P 


142 REACTIONS OF THE METALS 


REACTIONS OF CHROMIUM IN THE DRY WAY 


All chromium compounds color the borax, or salt of phosphorus, — 
bead an emerald green both in the oxidizing and reducing flames. 
Heated with sodium carbonate on charcoal before the blowpipe, all 
chromium compounds yield a green slag, which after long heating is 
changed to green infusible chromic oxide. By fusing with sodium 
carbonate and potassium nitrate in the loop of a platinum wire, all 
chromium compounds yield a yellow melt of alkali chromate: 


2Cre03+4Na2C03+302 = 4NaeCr04+4CO0e ; . 


If the fused mass is dissolved in water and acidified with acetic acid, 
the solution will give with silver nitrate a reddish-brown precipitate of 
silver chromate. This reaction is very delicate and serves for the 
detection of minute traces of chromium. Cloth which has been dyed 
with a chromium mordant can be tested in this way; the ash from a 
thread 5 cm, long is sufficient to give the test. 


IRON, Fe. At. Wt. 55.9 
Sp. Gr. 7.98. M. Pt. about 1600° C. 


Occurrence.—Native iron is rarely found. It occurs in basaltic 
rocks; also in meteorites, associated with nickel, cobalt, carbon, 
sulfur, and phosphorus. 

The most important iron ores are the oxides and the sulfides. 
Of these may be mentioned: 

Hematite, Fe2O3, isomorphous with corundum; magnetite, Fe30x, 
isomorphous with spinel; géthite, FeHO2, isomorphous with diaspore 
and manganite; limonite, Fe4H¢O9; (bog ore), Fe(OH)s, which is 
used in the purification of illuminating gas; pyrite, FeSe, which crystal- 
lizes in the isometric system; marcasite, FeSe, orthorhombic. Iron 
disulfide is, therefore, dimorphous. Another important iron ore is 
siderite, FeCO3, which is rhombohedral; vivianite, Fe3(PO4)2-8H20, 
is monoclinic. 

The metallic iron of commerce is never pure, but usually contains 
more or less iron carbide, iron sulfide, iron phosphide, iron silicide, 
corresponding manganese compounds and graphite, etc. 

On dissolving commercial iron in acids (H2SO4,HCl), hydrogen, 
contaminated with small amounts of hydrocarbons, hydrogen sulfide, 
mercaptans, phosphuretted hydrogen, and silicon hydride is given off, 
and these impurities give to the gas its unpleasant odor. There 
remains almost always an undissolved residue consisting chiefly of 
carbon. 


| 
4 
7 





TRON 143 


Tron is bivalent, trivalent, and rarely hexavalent, forming the 
following oxides: 


Iron protoxide Iron sesquioxide Ferrous-ferric oxide eg 
or ferrous oxide or ferric oxide or magnetite Iron trioxide 
FeO Fe203 Fe304 FeOs. 


Iron trioxide, FeO3, containing hexavalent iron, has never been 
isolated. It plays the part of an acid anhydride in ferrates of the general 
formula R2FeO4, which are decomposable by water. 

By dissolving these oxides in acid the corresponding salts are 
obtained; thus ferrous oxide gives with hydrochloric acid, ferrous 


: chloride, 


FeO0+2HCl=H20+ FeCl; 


ferric oxide gives ferric chloride, 
FeeO3 + 6HCl = 3HeO te 2FeCls ’ 


while ferrous-ferric oxide yields a mixture of ferrous and ferric chlorides: 
Fe304+8HCl =4H20+2FeCl3+FeCle. 


Iron, therefore, forms two-series of salts: first, the ferrous, derived 
from ferrous oxide, containing bivalent iron; second, the ferric, derived 
from ferric oxide, containing trivalent iron. These two series of salts 
show a quite different behavior toward reagents. 


The position of iron in the electromotive series shows that it is capable of 
being oxidized by hydrogen ions and suggests the possibility of its being 
oxidized to a measurable extent by the hydrogen ions of water. Careful. 
experiments have shown that absolutely pure water containing no dissolved 
oxygen, does not affect iron appreciably at ordinary laboratory temperatures, 
though there is evidence that a trace of iron dissolves and a film of hydrogen 
is formed on the metal which acts as a check upon further attack. The position 
of ferrous iron in the voltage series shows that hydrogen ions cannot oxidize 
iron appreciably to the ferric condition. The presence of dissolved oxygen, 
and this is normally present in all water that is exposed to the atmosphere, 
can accomplish this oxidation of the ferrous ions to the ferric condition and it 
also aids in the oxidation of the iron from the metallic to the ferrous state. In 
this way iron exposed to moisture and oxygen oxidizes or rusts. The rusting 
process is favored by the contact of the metal with a more noble metal, such as 
platinum, copper or nickel; an electric couple is formed and the iron becomes 
the positive pole, so that the hydrogen set free by the action of iron upon water 
is deposited upon the more noble metal. The presence of a less noble metal, 
such as zine, tends to hinder the corrosion of iron: the zinc corrodes instead of 
the iron. 

The presence of an acid is, therefore, not absolutely necessary to start the 
corrosion of iron. An increase in the concentration of hydrogen ions, however, 
will greatly hasten the solution of the metal. When carbonic acid is present, 
ferrous bicarbonate is first formed and, when the ferrous iron is oxidized to the 


144 REACTIONS OF THE METALS 
4 


ferric condition, the carbonic acid is set free again because ferric carbonate does 
not exist. The acid again acts upon the metal and the rate of its corrosion is — 
greatly accelerated. 

Certain substances tend to make iron passive, particularly strong nitric acid. 
Passive iron does not dissolve in dilute nitric acid and does not corrode readily 
(cf. aluminium, p. 126). On the other hand, certain substances can overcome 
the passive condition and are said to activate the iron. Thus a solution of com- 
mon salt is an activating agent. 

Different varieties of iron and steel corrode with different degrees of readi- 
ness. Cast iron is often protected by its casting skin. Impurities present in 
steel often favor corrosion by causing electric couples to be established, 


A. Ferrous Compounds > 


Ferrous compounds, which may be prepared by dissolving metallic 
iron, ferrous oxide, ferrous hydroxide, ferrous carbonate, or ferrous 
sulfide, ete., in acids, are usually greenish in the crystallized state, 
but in the anhydrous condition they are white, yellow or bluish; in 
concentrated solution they are green; in dilute solutions almost color- 
less. Ferrous compounds exhibit a strong tendency to change over into 
ferri¢ salts; they are strong reducing agents. 


‘REACTIONS IN THE WET WAY 


1. Ammonia produces in neutral solutions an incomplete pre- 
cipitation of white ferrous hydroxide: 


FeCle-+2NH3+2H20 @ Fe(OH)2+2NH4Cl. 


Ferrous salts in this respect are similar to those of magnesium 
(cf. p. 94). In the presence of ammonium chloride the reaction takes 
place in the direction from right to left; ammonia, therefore, causes 
no precipitation with ferrous salts out of contact with the air, provided 
sufficient ammonium chloride is present. On exposure to the air, 
however, a turbidity is soon formed, green at first, then almost black, 
and finally becoming brown. The small amount of ferrous hydroxide 
contained in the solution is oxidized by the air, forming at first black 
ferrous-ferric hydroxide and finally brown ferric hydroxide. 

2. Potassium and Sodium Hydroxides produce, if air is excluded, 
complete precipitation of white ferrous hydroxide, 


Fe+20H~ — Fe(OH)s2, 


which is quickly oxidized by the air into ferric hydroxide. 

3. Hydrogen Sulfide produces no precipitation in acid solutions 
of ferrous salts; in dilute neutral solutions a small amount of black 
ferrous sulfide is precipitated; but if the solution contains consid- 





IRON 145 


erable alkali acetate, hydrogen sulfide precipitates more of the iron as 
ferrous sulfide (but not all of it), in spite of the fact that ferrous 
sulfide is readily soluble in acetic acid. This interesting fact is an 
instructive illustration of the law of chemical mass action. 

The table on p. 22 states that 3.41078 -g. of ferrous sulfide dissolves 
in a liter of water. This small quantity exists in solution almost entirely as 
Fet* and S~ ions. FeS —@ Fet*++S>. When acetic acid, which is a much 
stronger acid than hydrogen sulfide (cf. p. 10), is added to the solution, equilib- 
rium has to be established between its hydrogen ions and the dissolved sulfide 
ions, 2H*+S~ — H.S, and, as a result of the formation of non-ionized hydrogen 
sulfide, the solution no longer contains enough sulfur ions to reach the value of 
the solubility product of FeS; to restore the equilibrium between FeS and its 
ions, more of the solid must dissolve. If, moreover, the solution is boiled, the 
hydrogen sulfide escapes as a gas as soon as it is formed. Consequently it is 
impossible to arrive at a state of equilibrium until all of the ferrous sulfide has 
dissolved. The solution is accomplished by means of hydrogen ions: 


FeS+2Ht = Fet+t+H,8 [. 


On the other hand, if the ionization of the acetic acid is repressed by adding 
an alkali acetate to the solution (cf. p. 46), and the concentration of the hydro- 
gen sulfide is made as large as possible by keeping the solution saturated with 
the gas, the reaction will take place in the reverse direction and some of the 
iron will be precipitated as ferrous sulfide. 

4. Ammonium Sulfide precipitates iron completely as black fer- 
rous sulfide: 

FeClo+(NH4)2S = 2NH4Cl+FeS, 


which is readily soluble in acids with evolution of hydrogen sulfide. 
In moist air it turns slightly brown, a part of the sulfur separates out, 
and a basic ferric sulfate is formed. 

5. Alkali Carbonates precipitate the white carbonate, 


FeCle +NaeCO3 =2NaCl +FeCOs, 
which in contact with the air becomes green, then brown: 
4FeCO; +6H20+02 =4CO>2 +4Fe(OH)s, 


being converted into ferric hydroxide with loss of carbonic anhy- 
dride. | 

Ferrous carbonate, like calcium carbonate (cf. p. 103), is soluble in 
carbonic acid, forming ferrous bicarbonate: 


FeCO: + HeCO3 = FeHo(CQOs) 2; 


a compound which is found in many natural waters, but which, like 
the normal carbonate, is decomposed by atmospheric oxygen with 
separation of ferric hydroxide: 


4eHs (COsz) 2 +2H:2O +02 = 8CO2 +4Fe(OH)s3. 


146 REACTIONS OF THE METALS 


Consequently a mineral water which contains ferrous bicarbonate, 
if allowed to stand in contact with the air, will become turbid, owing 
to the deposition of ferric hydroxide. To prevent this, the bottle 
must be filled with water and tightly corked, so that no trace of air 
can get in. Ferric hydroxide is insoluble in carbonic acid. 

6. Potassium Cyanide precipitates yellowish-brown ferrous cyanide, 


Fet++2NC” — Fe(CN)o, 


which is soluble in excess of the reagent, forming potassium ferro- 
cyanide: 
Fe(CN)2+4CN” — Fe(CN)6 ~. 


The complex ferrocyanide anion is in equilibrium, to be sure, with 
simple ferrous cations, but the quantity of the latter present in the 
aqueous solution of a ferrocyanide is so small that none of the above 
reactions characteristic of ferrous ions can be obtained with it. Many 
other similar complex cyanide anions are known; thus, the cyanides 
of silver, nickel, iron (ferrous and ferric), and cobalt all dissolve in potas- 
sium cyanide, forming the following: complex ions: [Ag(CN)e]", 
[Ni(CN)4]~, [Fe(CN).6] =, [Fe(CN)6]~~ [Co(CN)e]>. The acids are: 
H[{Ag(CN)2], He[Ni(CN)4], Hs[Fe(CN)6], HalFe(CN)6], Hs[Co(CN)sl. 
It is possible, as a matter of fact, to isolate the last three acids, though 
the two former have never been prepared; they immediately break 
down into metallic cyanide and hydrocyanic acid, just as carbonic acid 
is decomposed into water and carbon dioxide. 

With iron, therefore, there are two series of complex cyanogen 
compounds, the ferrocyanides and the ferricyanides. The ferro- 
cyanic derivatives contain the quadrivalent ferrocyanide anion and the 
ferricyanides contain the trivalent ferricyanide anion. 

Potassium ferrocyanide, K4{[Fe(CN).], is often called yellow prussiate 
of potash, and potassium ferricyanide, Ks3[Fe(CN)g¢|, is called red prus- 
siate of potash. The solubility of the alkali and alkaline-earth salts, and 
the insolubility and color of the salts of the heavy metals (especially 
with both ferric and ferrous iron), are very characteristic of ferro- and 
ferricyanides. 

7. Potassium Ferrocyanide, K,Fe(CN)», produces in solutions 
of ferrous salts, with complete exclusion of air, a white precipitate of 
potassium ferrous ferrocyanide or of ferrous ferrocyanide, depending 
upon whether one or two molecules of ferrous salt react with one mole- 
cule of potassium ferrocyanide: 


K4[Fe(CN)¢]-+FeSO4 = KoS04+KoFe[Fe(CN).] 
K,[Fe(CN)¢]+2FeSO4 = 2K2S04+Fe2[Fe(CN)¢] 


IRON : 147 


Although both of the above salts are white, a light-blue color is almost 
always obtained, because the precipitate is immediately oxidized some- 
what by the air, forming the ferric salt of hydroferrocyanic acid (Prus- 
sian blue): 

3Fe2[Fe(CN)¢6]-+-3H20+30 =Fe,[Fe(CN )elg3 +2Fe(OH)3. 

8. Potassium Ferricyanide, K3[Fe(CN).«], added to solutions of 
ferrous salts produces a dark blue precipitate (Turnbull’s blue) consist- 
ing of ferrous ferricyanide mixed with potassium-ferric ferrocyanide: 


2K3[Fe(CN).6]-+3FeCls =6KCl+Fes[Fe(CN).]2 


and 
Ks[Fe(CN).6]+FeCle = KFe|[Be (CN) 6]~~+2KCl. 


in other words the ferricyanide acts both as a precipitant and as an oxidizing 
agent * and a blue color results whenever iron is present in the cation in a state 
of oxidation different from that of the iron present in the complex anion. The 
ferricyanide ion is a strong oxidizing agent and in alkaline solution readily 
oxidizes ferrous hydroxide to ferric hydroxide. Turnbull’s blue is not very 
soluble in acid solutions, but is decomposed by treatment with caustic alkali, 
all of the complex anion being in the form of ferrocyanide: 


Fe;+ tiFe(CN)ele =4 SKOH = 2K,[Fe(CN)<| +2Fe(OH);+Fe(OH),, 
Kt¥Fet + t+[Fe(CN).] ==+3KOH =K,[Fe(CN).] +Fe(OH);. 


9. Potassium Thiocyanate gives no reaction with ferrous salts 
(note difference from ferric salts). 


As has been stated, ferrous salts are oxidized by the air to ferric salts; thus | 
ferrous sulfate is gradually changed into brown, basic ferric sulfate, 


2FeSO.+0 =Fe.0 -(SOx)s, 


’ which is insoluble in water. Consequently it often happens that ferrous sul- 
fate will not dissolve in water to a clear solution, but gives a brown, turbid 
solution, becoming clear on the addition of acid, the basic ferric salt being 
changed to a soluble neutral salt: 


Fe,0 : (SOx)2 +H.SO, == Fe, (SOx) 3+ H.0. 


Such a solution, which then contains ferric ions, reacts with potassium thio- 
cyanate (cf. p. 150). To free the solution from ferric salt, it may be boiled 
with metallic iron, with exclusion of air, whereby the ferric salt is changed into 
ferrous salt: 

Fe.(SO,);+Fe =3FeSO,. 


By means of strong oxidizing agents, ferrous salts can be quickly and 
completely changed into ferric salts, as was shown in the introduction (cf. 
pp. 27-33). 





* Cf. Erich Miuer, J. pr. Chem., 84 (1911), 353. 


148 _ REACTIONS OF THE METALS 


Detection of Ferrous Oxide in the Presence of Metallic Iron 


Treat the mixture with a large excess of a neutral solution of mercuric 
chloride and heat on the water-bath; the metallic iron goes into solution as 
ferrous chloride: 

2HgCl, 4. Fe= FeCl, + Hg,Cl. 


Filter off the residue and test the filtrate with potassium ferricyanide; a 
precipitate of Turnbull’s blue shows that metallic iron was originally present. 

Wash the residue with cold water, until all of the ferrous chloride has been 
dissolved, and then treat it with dilute hydrochloric acid. If the solution 
now gives a precipitate of Turnbull’s blue with potassium ferricyanide, ferrous 
oxide was present. ‘ 

If hydrogen is given off, some metallic iron is still present; the experiment 
must be repeated and the mixture given a longer treatment with HgCl, solution. 


B. Ferric Compounds 


Ferric oxide, Fe2O3, is reddish brown, becomes grayish black on 
strong ignition, but on being pulverized appears red again. 

The ferric salts are usually yellow or brown, but ferric am- 
monium alum is pale violet. Ferric salts are yellowish brown in aque- 
ous solution, and the solution reacts acid (hydrolysis). Dilution and 
warming favor the hydrolysis, so that all strongly diluted ferric salts 
deposit basic salts on being boiled: 


Fee(SO4)3+H20 @ Fe2(SO4)20+ H2S80.. 


With ferric salts of the weaker acids, often all of the iron is pre- 
cipitated as a basic salt; thus the acetate, on being boiled in a dilute 
solution, reacts as follows: 


Fe(C2H302)3-+2H20 @ Fe(OH) 2(C2H302) +2HC2H302. 


By the addition of acid all basic salts may be changed back into 
neutral salts. 


REACTIONS OF FERRIC SALTS IN THE WET WAY 
- 1. Ammonia precipitates brown, gelatinous ferric hydroxide: 
Fet*++30H- — Fe(OH)3. 


The solubility product of ferric hydroxide is so small (ef. p. 22) 
that it is precipitated completely even in the presence of ammonium 
salts; it is readily soluble in acids. On ignition it loses water and is 
changed to oxide, which is very difficultly soluble in dilute acids. It is 
best brought into solution by long-continued heating below the boiling 
point with concentrated hydrochloric acid. 


{RON 149 


2. Potassium and Sodium ee also precipitate ferric hy- 
droxide. 
3. Sodium Carbonate lice a brown precipitate of basic car- 


_ bonate, which at the boiling temperature is completely decomposed 


. > 


vs 


_hydrolytically into hydroxide and carbon dioxide: 


2FeCl3+3Na2CO3+3H20 =2Fe(OH)3+6NaCl+3C02 Tf. 


4, Zinc Oxide and Mercuric Oxide also precipitate the iron as 


hydroxide: 


2FeCls3+3Zn0+3H20 =3ZnCle+ 2Fe(OH)s. 


_ This reaction is frequently used in quantitative analysis. 


5. Sodium Phosphate precipitates yellowish-white ferric phosphate: 
FeCl3+2Na2HPO4 =3NaCl+ NaH2P01+FePOx,. 


Ferric phosphate is insoluble in acetic acid, but readily soluble 
in mineral acids. The precipitation of iron with sodium hydrogen 
phosphate is consequently only complete when a large excess of the 


‘precipitant is employed, or when sodium acetate.is added: 


FeCls +NasHPO, +NaCe2H30e =3NaCl + HC2H302 +FePO.. 


In this last case all the iron and all the phosphoric acid are pre- 
cipitated. The reaction is often used to precipitate phosphoric acid 
quantitatively. An excess of the disodium phosphate will also cause — 
complete precipitation of iron as phosphate, if the phosphate solution 
is previously exactly neutralized with ammonia: 


NazHPO4+NH40H = H20+NazNH4POu, 
and 
NaeNH4P04+ FeCls = 2NaCl+ NH4Cl+FePO.. 


If, however, an excess of sodium phosphate and ammonia is added 
to the iron solution, the precipitation of iron is incomplete, because the 
ferric phosphate dissolves in the excess of sodium phosphate, in the 
presence of ammonia (or ammonium carbonate), with a brown color 
and formation of a complex salt. 

Ferric phosphate is transformed by ammonia into a brown basic 
phosphate, and by potassium hydroxide almost completely into ferric 
hydroxide and potassium phosphate; while by fusion with caustic 
alkali or alkali carbonate it is completely decomposed. 

If alkaline earth ions are present, an excess of ammonia completely 
changes ferric phosphate to ferric hydroxide and alkaline carth phos- 
phate is precipitated. 


150 REACTIONS OF THE METALS 


6. Alkali Acetates produce in cold, neutral solutions a dark-brown 
coloration, and on boiling the dilute solution all of the iron separates 
as basic acetate: 


FeCl3+3NaC2H302=3NaCl+Fe(C2H302)3 (in the cold), - 
Fe(C2H302)3-+2H20 = 2HC2H302+Fe(OH)2C2H30z (on boiling). 


The presence of organic hydroxy-acids (tartaric, malic, citric, ete.) 
and of polyatomic alcohols (glycerol, erythritol, mannitol, sugars, etc.) 
prevent all of the above-mentioned reactions, because complex salts 
are formed in which the iron is present in the form of a complex anion 


(cf. aluminium, p. 128). 
‘7. Potassium Thiocyanate, KCNS, produces in solutions of ferric 


salts a blood-red coloration: 
Fet+++3CNS~ = Fe(CNS)s. 


This action is reversible; the red color of the slightly ionized ferric thio- 
cyanate being most intense when an excess of ferric salt, or of potassium thio- 
cyanate is present. 

If the solution is shaken with ether, the Fe(CNS); goes into the ether. Ferric 
thiocyanate combines readily with potassium thiocyanate, forming complex 
potassium ferrithiocyanate: 


Fe(CNS);+3KCNS =K,[Fe(CNS)].* 


analogous to potassium ferricyanide, K,[Fe(CN)s]. 

The complex salt is insoluble in ether, the Fe(CNS); only being soluble 
therein; so that the red color is due to the formation of the ferric thiocyanate 
and not to the complex salt. 

This reaction is extremely sensitive, but cannot always be relied on. The 
test cannot be made in the presence of strong oxidizing agents such as nitric 
acid, as a red color is produced by the oxidation of the thiocyanate. The 
oxidized compound is not very stable, however, and its color is not very deep. 
If the solution contains considerable alkali acetate, the coloration cannot be 
recognized. The presence of organic hydroxy-compounds (tartaric acid, ete.) 
prevents the reaction in neutral solutions, but not in acid solutions. In the 
presence of mercuric chloride the red color disappears entirely; the mercuric 
chloride reacts with the ferric thiocyanate, forming a colorless, soluble 
mercuric double salt, which is ionized even less than ferric thiocyanate: 


2Fe(CNS);+6HgCl, =2FeCl,;+3[Hg(CNS).-HgCl.]. 
8. Potassium Ferrocyanide, K4Fe(CN).6, produces in neutral or 
acid solutions of ferric salts an intense blue precipitation of Prussian 


blue: : 
3[Fe (CN) 6] ee 4Fet tt, Fe, [Fe (CN) 6ls- 





* K;[Fe(CNS)«]+4H.0. Cf. RosenueErm, Z. anorg. Chem., 27 (1901), 208. 


IRON 151 


Prussian blue, the ferric salt of ferrocyanic acid, is insoluble in water, but 
soluble in oxalic acid and in an excess of potassium ferrocyanide; the solution 
thus obtained is a deep blue and is used as blueing and as blue ink. The blue 
solution obtained with a ferric salt and an excess of potassium ferrocyanide con- 
tains colloidal KFe[Fe(CN).]-H.O, which can be salted out by the addition of a 
considerable quantity of electrolyte such as alkali chloride, sulfate or nitrate. 
Prussian .blue is also soluble in concentrated hydrochloric acid, but is pre- 
cipitated again on dilution. As the ferric salt of ferrocyanic acid it behaves 
like other ferric salts to the hydroxides of the alkalies; ferric hydroxide and the 
alkali salt of hydroferrocyanic acid being formed: 


Fe,[Fe(CN).];+120H- — 4Fe(OH);+3[Fe(CN),|=. 


9. Potassium Ferricyanide, Ks|Fe(CN).»], produces no precipita- 
tion in solutions of ferric salts, only a brown coloration (differing 
from ferrous salts) : 


[Fe(CN).6]=+Fet ** = 2Fe(CN)s3. 


10. Ammonium Sulfide, added to a solution of a ferric salt, gives 
a precipitate of ferric sulfide, Fe2Ss, 


2Fet $y T439= <> Fe2Ssz, 


which is soluble in cold, dilute hydrochloric acid, forming ferrous 
chloride and sulfur: 


Fe.S3+4H+t > 2Fet++2H.S fT +S. 


The fact that Fe.S; is precipitated, and not FeS as commonly believed, was 
proved by H. N. Stokes* who decomposed it out of contact with air by zinc- 
ammonium oxide and obtained white ZnS and red Fe(OH);._ L. Gedelt has also 
shown that hydrogen sulfide passed into a solution of ferric chloride made 
alkaline with ammonia gives Fe.8,. If, however, the solution is acid, hydrogen 
sulfide or ammonium sulfide reduces the iron before any precipitate is formed. 


11. Hydrogen Sulfide in acid solutions reduces ferric salts to 
ferrous salts, with separation of sulfur: 


2Fe*++t++H.S — 2Fet*+2Ht-+S. 


Besides hydrogen sulfide, many other substances (nascent hydrogen, 
stannous chloride, sulfurous acid, hydriodic acid, etc.) will reduce ferric 
salts, as was shown on pp. 35, 36. 

12. Ether when shaken with a solution of ferric chloride in 6 N 
hydrochloric acid dissolves most of the ferric chloride. By separating 





* J. am. Chem. Soc., 29 (1907), 304. 
t Ueber Schwefeleisen, Karlsruhe, (1905). 


152 REACTIONS OF THE METALS 


the ether with the aid of a separatory funnel, and repeating the opera- 
tion, nearly all of the iron can be removed from the aqueous solution. 
(CH, 7... 1%) 

13. Cupferron, the ammonium salt of phenylnitrosohydroxylamine, 
CesHsNO-NONHg, precipitates red (CeHsNO-NO)s3Fe, which is soluble 
in ether, insoluble in acids, and converted into Fe(OH)3 by treatment 
with ammonia. 

14. Sodium Thiosulfate, NazSeOQ3, colors neutral ferric solu- 
tions a violet red, but the color disappears quickly and the solution 
then contains ferrous salt and sodium tetrathionate; 


2Nao8003+-2FeCls = 2NaCl+2FeClo+NaeSi0c. 


The composition of the violet-red substance which is first formed 
is unknown; perhaps it is ferric thiosulfate. 


As we have seen, there exist a number of iron compounds which contain 
the metal as a complex ion, so that it cannot be detected by the ordinary 
reagents. The complex hydroxy-organic compounds, as well as the ferro- 
and ferricyanide compounds, belong to this class of compounds. 

If it is a question of proving the presence of iron in such a compound, a 
different method should be used in the case of an organic hydroxy-compound 
from that in the case of a ferro- or ferricyanide. 

If organic substances are present, the iron is precipitated as sulfide by means 
of ammonium sulfide; or the organic matter is first removed by ignition, 
whereby metallic iron, ‘oxides of iron and carbon are obtained. 

In case we have a ferro- or ferricyanide, the iron cannot even be piecinieeae 
by means of ammonium sulfide; the compound must be completely destroyed 
before it will be possible to detect the peer? of iron by any of the ordinary 
methods. 

This may be accomplished (a) by ignition, (6) by fusion with potassium | 
carbonate or sodium carbonate, or (c) by heating strongly with concentrated 
sulfuric acid. 

(a) Decomposition by Ignition.—The ferrocyanides are decomposed (with 
evolution of nitrogen) into potassium cyanide and carbide of iron; 


K,[Fe(CN). =4KCN+FeC.+N, tT . 


The ferricyanides also leave behind iron carbide and potassium cyanide, 
but evolve cyanogen as well as nitrogen: 


2K;[Fe(CN).] =6KCN+2FeC,+(CN).+2N, fT. 


Treat the residue from the ignition with water, whereby the potassium 
cyanide goes into solution, leaving behind the iron carbide; filter and treat 
- the residue with hydrochloric acid. The iron goes into solution as ferrous 
chloride, hydrocarbons are given off, and there remains some carbon. 

The above decomposition can be imagined to take place as follows: 


IRON 153 


By heating potassium ferrocyanide, it is decomposed first into potassium 
cyanide and ferrous cyanide, while the latter on further heating is changed 
to iron carbide and nitrogen: 


(a) K.[Fe(CN)«] =4KCN +Fe(CN):; 
(8) Fe(CN)2=FeC:+N, T. 


Potassium ferricyanide is decomposed into potassium cyanide and the 
- very unstable ferric cyanide, which splits off cyanogen and becomes ferrous 
cyanide; the latter is decomposed, as before, into iron carbide and nitrogen: 


(a) K;[Fe(CN).] =3KCN +Fe(CN);; 
(8) Fe(CN);=Fe(CN).+CN T ; 
(y) Fe(CN). =FeC, +N, t ° 


(b) Decomposition by Means of Fusion with Potassium Carbonate.—Mix the 
substance with an equal amount of the carbonate and heat in a porcelain cru- 
cible until a quiet fusion is obtained. By this means a mixture of potassium 
cyanide and potassium cyanate (both soluble in water) is formed in the presence 
of metallic iron: 


K,[Fe(CN).]+K.CO; =5KCN+KCNO+CO, f +Fe. 


Extract the melt with water, filter and dissolve the iron in hydrochloric 
acid. 

(c) Decomposition by Heating with Concentrated Sulfuric Acid.—By heating 
with concentrated sulfuric acid all complex cyanogen compounds may be 
decomposei. By this means the metal present is changed into sulfate, the 
nitrogen of the cyanogen into ammonium sulfate, while the carbon of the 
_ cyanogen escapes as carbon monoxide: 


K,[Fe(CN).] +6H,S0,+6H,0 =2K,S0,+ FeSO,+3(NH,).S0,+6CO Tf , 
2K;[Fe(CN).] +12H,SO,+12H,0 =3K.80,+ Fe.(SO,); +6(NH,)2S0,+12CO Tf . 


The treatment with concentrated sulfuric acid is best accomplished in a 
procelain crucible placed in an inclined position over the flame, and the flame 
directed against the upper part of the crucible. Continue heating until fumes 
of sulfuric acid cease to come off. Treat the residue, which consists of an 
alkali sulfate and anhydrous ferrous or ferric sulfate, with a little concen-— 
trated sulfuric acid, heat gently, and add water little by little. In this way the 
sulfate is readily brought into solution. 


REACTIONS IN THE DRY WAY > 


The borax (or sodium metaphosphate) bead, containing a small 
amount of an iron salt, is yellow while hot and colorless when cold 
after being heated in the oxidizing flame, and pale green after being 
heated in the reducing flame. When strongly saturated, however, the 
bead obtained with the oxidizing flame is brown while hot, yellow 


» \ 


154 REACTIONS OF THE METALS 


when cold; and after heating in the reducing flame it becomes 
bottle-green. 

Heated on charcoal with soda before the blowpipe, all iron com-: 
pounds leave a gray particle of metallic iron, which is usually difficult 
to see, but can be separated from the charcoal by means of the magnet. 
The reduction on the charcoal stick, as described on p. 65, is a 
much more delicate test. 


URANIUM, Ur. At. Wt. 228.5 
Sp. Gr. =18.33. M. Pt.=2500(°?) 


Occurrence.—Uranium occurs in nature chiefly in the mineral 
pitch-blende, U30g; but it is also found in a few rare minerals, uranite, 
(UO2)2Cu(PO4)2+8H20; samarskite (a niobite of iron, yttrium, 
cerium, and erbium with varying amounts of uranium); and liebigite, 
U(COs)2-2CaCO3+10H20. 

Klaproth showed, in 1789, that the mineral pitch-blende contained 
a new metal, which he called uranium. By heating the oxide with 
reducing agents he obtained a brown, almost copper-red, substance, © 
which he took to be the metal, and it indeed does behave like a metal, 
dissolving in acids in contact with the air, forming yellowish-green salts. 

It was not until 1842 that it was shown by Péligot that this reddish- 
brown body was not the metal uranium, but its dioxide. The hexa- 
valent metal itself was obtained by Péligot, as a gray powder, by reduc- 
ing the tetrachloride with sodium. 
| Out of contact with the air, uranium dioxide (uranyl) dissolves in 

strong acids, forming wranous salts: 


U02+4HCl=2H20+UCh, UO2+2H2804 = 2H20+ U(SO4)o. 


The uranous salts are extremely unstable, and on being exposed 
to the air change rapidly, forming wranyl salts which contain the 
bivalent UOz group: 


UClk,+0+H20 — UOeCle+2HCl, 
U(SO4)2+0+He20 = U02804+ HeSOu. 


Only the reactions of the uranyl salts will be described in this book. 

Besides uranyl (or uranium dioxide) uranium forms a trioxide, 
UOz, which can be regarded as uranyl oxide, UO2O. It dissolves 
in acids, forming urany] salts: 


U03+2HCl = UOe2Cle+H20. 


- URANIUM 155 


By igniting the oxides of uranium in air, dark-green urano-uranic 
oxide, U3Og or (2U03-UQOz2), is obtained, which out of contact with 
the air dissolves in strong acids, forming a mixture of uranous and 
uranyl] salts: 


(2U03-UO2) +4H2804 = 2U02804+ U (S04)2+4H20. 
By dissolving in aqua regia, uranyl chloride is obtained: 
3U30s+18HCl1+2HNO3=9U02Cl2+2NO 7 +10H20. 


All uranyl compounds are colored yellow or yellowish green. Most 
of them are soluble in water, but the oxides, the sulfide, phosphate, 
and uranates are insoluble. In mineral acids all uranium compounds 
are soluble, with the exception of the ferrocyanide. 


REACTIONS OF URANYL COMPOUNDS IN THE WET WAY 


1. Potassium Hydroxide precipitates yellow amorphous potassium 
uranate. Uranyl hydroxide, UO2(OH)s, is first formed and changes 
into uranic acid, H2U207, of which the alkali salts are insoluble: 


2U02Clz+6KOH — K2U207+4KC1+3H20. 
2. Ammonia precipitates yellow, amorphous ammonium uranate: 
2U02 (NOs) 2 +6NH40OH os (NH4)2U207 +4NH4NOs +3H20. 


The alkali uranates are soluble in alkali carbonates, particularly in ammo- 
nium carbonate, with the formation of complex salts: 


(NH,).U20;+6(NH,4)2CO;+3H20 — 2(N H,).{UO.(CO;)3]++6NH,OH. 


Consequently, in the presence of sufficient alkali carbonate, ammonia fails to 
precipitate uranium. ‘Tartaric and citric acids (and other organic substances) 
also prevent the precipitation with ammonia and caustic alkalies, as with iron, 
chromium, and aluminium. — 


3. Sodium Carbonate produces in concentrated solutions an orange- 
yellow precipitate of sodium uranyl carbonate: 


UO2(NO3)2+3Na2CO3 = 2NaN03+Na4[U02(CO3)3]. 


Sodium uranyl carbonate is soluble in considerable water, so that 
no precipitate is formed from dilute solutions. It is still more soluble 
in alkali carbonate solution, particularly in a bicarbonate solution. 
From such solutions sodium hydroxide precipitates sodium uranate, 
but ammonia produces no precipitation. 


156 REACTIONS OF THE METALS 


4, Barium Carbonate precipitates in the cold all of the uranium, ral 
probably as barium uranyl carbonate: 


UOe2(NO3)2+8BaCO3 = Ba(NOz)2+Ba2{U02(COs3)3]. 
5. Ammonium Sulfide precipitates brown uranyl sulfide, 
UO2(NOs)2+(NH4)2S = 2NH4NO3s+U0O2S, 


soluble in dilute acids and in ammonium carbonate: 


U028+3(NH4)2CO3 = (NH4)28+ (NH4)4[U02(COs)3]. 


Ammonium sulfide, therefore, produces no precipitate in solutions of 
uranyl salts in the presence of ammonium carbonate. 
6. Sodium Phosphate precipitates yellowish-white uranyl phosphate, 


UOe2 (NOs) a+ NazgHP O4 —- 2NaNOsz +UO2HPO,, 


while, in the presence of ammonium acetate, uranyl ammonium phos- 
phate is precipitated: 


NazgHPO, +U0e2 (NOsz)2 +NH4CoH302 ae 
2NaNOz +HCe2H302 +U02 . NH,4POx,. 


Both precipitates are insoluble in acetic acid, but soluble in mineral 
acids. 

7. Potassium Ferrocyanide produces a brown precipitate, or in 
very dilute solutions, a brownish-red coloration, 


Fe(CN)¢ ~+2U02t* — (UO2)2[Fe(CN)<], 


On addition of potassium hydroxide the brownish-red_precipi- ! 
tate becomes yellow, owing to the formation of potassium uranate: 


(UO2)2[Fe(CN)6]+6KOH = Ka[Fe(CN)¢6]+3H20+K2U207. 


_ (Distinction from cupric ferrocyanide.) 


REACTIONS IN THE DRY WAY 


The borax (or sodium metaphosphate) bead is yellow in the oxidizing — 
flame and green in the reducing flame. 


TITANIUM 157 


TITANIUM, Ti. At. Wt. 48.1 
Sp. Gr. =4.87. M. Pt. =1900° (?) 


Occurrence.—Titanium occurs in nature most frequently as the 
dioxide, rutile (tetragonal), anatase (tetragonal), and brookite (ortho- 
rhombic). Titanium is also found in the minerals perowskit, CaTiOs, 
titanite, CaSiTiOs, and ilmenite, FeTiO3, as well as in many crystalline 
rocks. It is present in most rocks, but usually only in very small 
quantity. 

Titanium itself is a gray metal, very similar to iron. On being 
heated in the air it burns brightly to white titanium oxide. The 
following oxides of titanium are known: TigQ2, TizO3, TiOe, TiO3 or 
possibly TiO2-H2QOz. 

The oxides TigO2 and TigO3 form violet-colored salts, which are 
readily changed by oxidizing agents into derivatives of TiOz. The 
most important oxide is titanium dioxide, which sometimes acts as 
a base and sometimes as an acid. The titanium dioxide as it occurs 
in nature (rutile, etc.) is insoluble in all acids. In order to bring it 
into solution it is best to fuse it with potassium pyrosulfate, whereby 
it is changed into titanium sulfate: 


Ti02+2K28207 = Ti(SO4)2+2K280u. 


The melt is dissolved in cold water. It can also be dissolved by fusing 
with sodium carbonate and treating the melt with 6 N hydrochloric 
acid. 

REACTIONS IN THE WET WAY 


For these reactions a solution of titanium sulfate or of titanium 
hydroxide in hydrochloric acid may be used. 

1. Potassium Hydroxide precipitates, in the cold, gelatinous 
orthotitanic acid, 


++ 
Tit++40H- > HuTiO,, 


which is almost insoluble in an excess of the reagent, but readily soluble 
in mineral acids. 

If the precipitation by potassium hydroxide takes place from a 
hot solution the titanium is precipitated as metdtitanic acid, 


++ 
Tit ++40H — H20+H2TiOs, 


which is difficultly soluble in dilute acids. By long digestion with 
concentrated hydrochloric or sulfuric acid it goes gradually into 
solution. By the ignition of both these titanic acids the anhydride 


158 REACTIONS OF THE METALS 


TiO>s is obtained, which is only slightly soluble in concentrated hydro- 
chloric acid, but readily soluble in hot concentrated sulfuric acid. 

2. Ammonia, Ammonium Sulfide, and Barium Carbonate (like 
potassium hydroxide) precipitate, in the cold, orthotitanic acid, readily 
soluble in acids; and from hot solutions the difficultly soluble meta- 


titanic acid. 
TiCl4+4N H40OH =4NH4Cl+ HygTiOg, 


TiCla+ 2(NH4)2S+4H20 =4NH4Cl+2H2S 7 +HyTiO,, 
TiCl,+2BaCO3+2H20 =2BaCle +2COs T +HuTiOs. 


3. Alkali Acetates precipitate on boiling all of the titanium as 
metatitanic acid: 


++ 
Ta +44CoH302 +3H20 =4HC2H302+ HeTiOs3. 


Titanium acetate is first formed, but it is completely decomposed 
hydrolytically by boiling the dilute solution. 

4. Water.—Not only titanium acetate is hydrolytically decom- 
posed by water, but all titanium salts. | 


The ease with which soluble titanium salts undergo hydrolysis with the 
formation of insoluble titanic acid is the basis of the several methods for 
separating titanium from aluminium, iron, chromium, ete.; the oxides of these 
metals are fused with potassium pyrosulfate, the product of the fusion is 
dissolved in cold water,* and the solution is then heated to boiling. The tita- 
nium is completely precipitated as granular metatitanic acid, which can be’ 
readily filtered off, while the remaining metals remain in solution as sulfates: 


Ti(SO,) +3H.0 = 2H.80,+-H.TiO;. 


As this reaction, like all hydrolytic decompositions, is reversible, it is evident 
that, in order to make the precipitation of the metatitanic acid complete, the 
amount of free acid present should be kept small,{ considerable water should be 
used, and the solution kept hot while filtering. 

In order to precipitate titanic acid from a solution according to this method, 
the concentrated solution it treated with sodium carbonate in the cold until 
a slight permanent precipitate of Ti(OH), is obtained, sulfuric acid is added 
drop by drop until the precipitate is just dissolved, considerable water is added 
(300 to 500 cc. of water should be used for each 0.1 gm. TiO.) and the solution 
kept at the boiling temperature for one hour. The granular metatitanic acid 
thus obtained is easy to filter as long as free acid is present. On being washed 
with pure water a turbid filtrate is always obtained; a little dilute sulphuric 
acid should therefore be added to the wash-water. 


* Solution takes place much more quickly if the liquid is kept in constant motion, 
e.g., by conducting a current of air through it. 
+ If too little acid is present, however, iron and aluminium will precipitate. 





TITANIUM 159 


f 


With the separation of titanium according to the above-described method, 
a play of colors will be observed in the bottom/of the glass beaker or flask which 
is very characteristic of titanic acid. 

The presence of tartaric acid, citric acid, and many other organic compounds 
prevents the above reactions. ‘In such a case the organic substance must be 
first destroyed either by ignition or by oxidation with potassium permanganate 
(see pp. 129 and 152), the titanium dioxide dissolved in sulfuric acid and pre- 
cipitated according to any of the above methods, 


5. Potassium Ferrocyanide produces in slightly acid solutions 
a brown precipitate. 

6. Tannin produces a brown precipitate, which soon becomes 
orange. 

7. Sodium Thiosulfate precipitates in boiling solutions all of the 
titanium as metatitanic acid: 


TiCk,+ 2Na25203-+ H20 =4NaCl+2S+2802 T +HoeTiOs. 
8. Sodium Phosphate precipitates basic titanium phosphate, 


TiCl4a+3Na2HPO4s.+ H20 =4NaCl+2NaH2P04+Ti‘PO,) (OH), 


soluble in mineral acids, insoluble in acetic acid. 

9. Hydrogen Peroxide.—If hydrogen peroxide is added to a slightly 
acid solution of titanium sulfate, the solution is colored orange red, 
except in the presence of small amounts of titanium, when the color is 
light yellow. This reaction, which depends upon the formation of 
TiO3, or possibly TiO2-H2QOs2, is exceedingly delicate, and is especially 
suitable for the detection of titanium in rocks. Vanadic acid behaves 
similarly with hydrogen peroxide. 

If a solution of titanium sulfate is treated with a large excess of 
hydrogen peroxide and then potassium hydroxide is added, a precipitate 
is formed which dissolves in a great excess of the alkali, forming a 
yellow solution. This solution remains clear for a long time, but 
eventually a bright yellow precipitate of Ti(OH)4 is formed.* 

An insoluble titanium compound on being fused with sodium 
peroxide in a nickel crucible yields a melt which permits the extraction 
of all the titanium by water.+ If the solution is made strongly acid 
with sulfuric acid, the orange-red color of pertitanic acid is apparent. 
If iron was present in the original insoluble titanium mineral, it is left 
insoluble in water after the fusion with sodium peroxide, 





* A. CLAssEN, Ber., 21 (1888). 
tJ. H. Watton, Jr., J. Am. Chem. Soc. (1907). 


160 REACTIONS OF THE METALS 
10. Zinc or Tin .produces in acid solutions, preferably hydro- 
chloric acid, a violet color caused by the formation of TigCle: 


2TiCl4 -}- He = 2HCl + TisCle. 


The quadrivalent titanium compounds are not reduced by hydrogen 
sulfide or sulfurous acid. 

11. The Fluoride is quantitatively éhanged to the dioxide by 
evaporation with sulfuric acid (difference from silicic acid). 


TiF1+2H2S04=4HF T +2803 T +TiOo. 


12. Ether does not dissolve titanium chloride. By shaking the 
hydrochloric acid solution with ether, therefore, it is possible to separate 
ferric chloride from titanium chloride. 

13. Cupferron cf. (p. 152) precipitates the yellow titanium salt of 
phenylnitrosohydroxylamine from acid solutions: 


++ 
Tit *+4CsH;sNO-NO™~ —> Ti(CsHs;NO-NO)s. 


In the presence of tartaric acid, iron can be removed by ammonia 
and ammonium sulfide and the titanium precipitated in the acidified 
filtrate with cupferron. 


REACTIONS IN THE DRY WAY 


Titanium compounds do not color the borax or sodium meta- 
phosphate bead in the oxidizing flame; after continued heating in’ 
the reducing flame the bead becomes yellow while hot and violet 
when cold. By the addition of a little tin the violet color appears 
much more quickly. The addition of iron causes a brownish to red 
bead. 

On fusing titanic acid with sodium carbonate, sodium metatitanate 
is formed, which is readily soluble in acids. By treatment with hot 
water, sodium metatitanate is decomposed, forming metatitanic acid, 
which is difficultly soluble in dilute acids. Even in cold water hydrolysis 
of sodium titanate takes place and none of the titanium dissolves, 


MANGANESE, Mn. At. Wt. 54.93 
Sp. Gr. =about 8.0. M. Pt. =1225° C. 


Occurrence.—The most important manganese minerals are pyrolu- 
site, MnOzg, orthorhombic; polianite, also MnOog, tetragonal, isomor- 
phous with rutile and tinstone; braunite, Mn2Os, tetragonal; man- 
ganite, HMnOz, orthorhombic, isomorphous with géthite and diaspore; 
hausmannite, MngQO4, tetragonal; and rhodochrosite, MnCO3s. Man- 


* MANGANESE 161 


ganese is a constant companion of iron, so that we find it in varying 
amounts in almost all iron ores. 

It is ‘a grayish-white metal which is readily oxidized in moist air, 
and is attacked by dilute acids, even acetic acid. It forms the follow- 
ing oxides: 

MnO, Mn2Qs, Mns3QO0u4, MnOz, (MnQOs), Mne207. 

By treating any one of these oxides, except MnO, with cold, dilute 

hydrochloric acid a dark, greenish-brown solution is obtained, which 


on. being heated evolves chlorine and becomes colorless. The solution 
then contains a bivalent manganese salt,—a derivative of MnO. 


Mn0O+2HCl=H20+MnCle, 

Mn203 +6HCl=3H20 +2MnCle+Cle T , 
Mn304 +8HCl=4H20 +3MnCly+Cle T , 
Mn0O2+4HCl=2H20 +MnCly+Clo T , 
Mn:207 +14HCl=7H20 +2MnCle+5Cle 7. 


All manganese oxides dissolve on warming with concentrated 
sulfuric acid, forming manganous sulfate, accompanied (except with 
MnO) by evolution of oxygen: 


MnO +H2S04=H20 +MnS80,, 

2Mn203 + 4H2SO4 =4H20 +4Mn804+02 Tf, 
2Mn304 +6H2801 =6H20 +6MnS04+02 T, 
2MnOz +2H 2804 =2H20 +2MnS04+02 1. 


The behavior of the higher oxides, MnOz, MneQ3, and Mn3Q0q, 
with boiling dilute nitric or sulfuric acid is very interesting: MnOs 
is not attacked at all by these dilute acids; while Mn2Oz3 gives up 
half of its manganese to the acid, the other half remaining undissolved 
as brown hydrated manganese dioxide, H2MnOs; two-thirds of Mn304 
is dissolved by these acids, brown hydrated manganese dioxide being 
left behind, as before. 

This HeMnO3 separates out just as metasilicic acid is deposited 
from a silicate on the addition of a strong acid: 


CaSiOz + QHt > HeSi03 + Cat*. 


In fact, hydrated manganese dioxide behaves in most cases exactly 
like an acid, the oxides Mn2Oz and MngQOx behave like manganous salts 
of this acid and are to be regarded as manganites. 


162 REACTIONS OF THE METALS 


Mn203 therefore is to be regarded as manganous manganite, 
O Ne 
Mn 0 MnO, of analogous composition to manganous carbonate, 


MnCOs, and manganous metasilicate, MnSiO3. 

According to this conception, electrolytic dissociation shou!d give 
rise to Mn** cations and MnO; anions; it is easy to understand, there- 
fore, why Mne2O3 gives up half of its manganese on treatment with 
dilute nitric acid, with the separation of manganous acid: 


Mn:-Mn0O;+2HNO;3 @ Mn(NOs3)2+H2MnOs. 


Mn3zQ4, which gives up two-thirds of its manganese, may be con- 
sidered to be the manganous salt of orthomanganous acid, Hy4MnQ,. 

On treating Mnz-MnO, with nitric acid, the ortho acid first sep- 
arates out; it loses water, and goes over into metamanganous acid: 


Mn2-Mn0,+4HNO;3 = 2Mn(NOs3) 2+H4Mn0Oxz, 
H4zMn0O, == HO +H2MnOs3. 


MnOz stands in the same relation to HoMnOs as COz to HeCOs, 
as SiO2 to HeSiOs, and as SnOzg to HeSnOs; MnOsz, therefore, behaves 
like an acid anhydride. 

Like SnO2g (which see), manganese dioxide behaves partly as an 
acid anhydride and partly as the anhydride of a base, because it prob- 
ably forms the chloride MnCl4. For if MnOz is treated with cold 
concentrated hydrochloric acid it dissolves with a brownish-green 
color, forming manganese tetrachloride, soluble in ether with a green 
color. If, therefore, the aqueous solution of MnCl, is shaken with 
ether, the upper layer is colored green. 


Mn(S0,)2 and Mn,.(SO,); are also known. These salts are hydrolyzed 
readily, forming H,2MnO; and Mn-Mn0O;: 


Mn (SOx)2 +3H,0 = H,Mn0O; +2H.SO,, 
Mn,.(SO,);+3H.0 @ Mn,0;+3H.S8O.. 


Not only manganous manganites are known, but quite a number 
of other manganites. Some of these play a very important part in 
analytical chemistry; as, for example, zine and calcium bimanganites, 
ZnHe(MnOs)2, CaH2(MnOs)2, which are analogous to calcium bicar- 
bonate, CaH2(COs)e. 

Zinc bimanganite is formed in the volumetric determination of 
manganese (see Vol. II). Calcium bimanganite is of importance. 
technically. Thus the recovery of manganese, by the Weldon process, 


MANGANESE 163 


in the manufacture of chlorine, depends upon the formation of calcium 
bimanganite. 

Manganous oxide, MnO, is the only oxide of manganese which 
in all cases acts as the anhydride of a base. By dissolving this oxide 
in acids, manganous salts are always obtained, in which the manganese 
is bivalent. ‘The oxide MnQOg has never been isolated, but there are 
salts (ReMnOuz, see p. 169) known which are derived from it. Mng207 
is a distinct acid anhydride, from which the permanganates (RMnQ,) 


are derived. 
In the study of the reactions of manganese we will consider first 
the manganous compounds, then the manganates and permanganates. 


A. Manganous Compounds 


The manganous compounds are pink both in the crystalline state 
and in aqueous solution; but in the anhydrous state they are colorless 
with the exception of the sulphide. 


REACTIONS IN THE WET WAY 


1. Potassium or Sodium Hydroxide precipitates white manganous 


hydroxide, 
Mn*tt+20H- — Mn(OH)p, 


which rapidly becomes brown in the air, owing to the formation of 
manganous manganites, which are less soluble than Mn(OH)e. 
First a part of the manganous hydroxide is oxidized by the air 


- to manganous acid: 
Mn(OH)2+0 — H2Mn0Os, 


which, on coming in contact with the basic manganous hydroxide, 
immediately forms a salt with it—a manganite, 
H2Mn0O3+Mn(OH)e2 <2 Mn-Mn0O3+2H20 
or possibly, 
2HeMnO3 +Mn(OH)s2 = MnHo2(Mn0Osz)e +2H>0. 
This oxidation takes place in the air only gradually, but imme- 


diately in the presence of chlorine, bromine, hypochlorites, hydrogen 
peroxide, etc.: 


Mn(OH)2+2Na0H+ Cle = 2NaCl+ H20 +H2MnOs, 
Mn(OH)2+H202 = H20+H2Mn0O3. 


164 REACTIONS OF THE METALS 


The formation of manganites is of technical importance, as mentioned on 
p. 162. The residue obtained in the preparation of chlorine from pyrolusite 
and hydrochloric acid consists chiefly of manganese chloride; by adding lime 
to it, manganous hydroxide is formed. This mixture of manganous hydroxide 
and lime is exposed to the action of the atmosphere, whereby manganous acid 
is formed, which unites with the calcium as the stronger base, forming calcium 
bimanganite, so that finally all of the manganese is oxidized to manganous acid: 


2Mn(OH), +0, +Ca(OH), = 2H.0 +CaH,(Mn0O;)>. 


On treating the residue, when in the right condition. with hydrochloric 
acid again, the same amount of chlorine is obtained as from the original pyro- 
lusite: 

2Mn0O.+S8HCl = 4H,.0+2MnCl, +2Cl, T 4 


CaH.(MnO;),+10HCl =6H,0+2MnCl,+CaCl,+2Cl, T. 


It is, however, necessary to add a little more hydrochloric acid in the latter 
case, because a part of the acid is used up in setting the An gSEos acid free 
from the manganite. 


2. Ammonia precipitates (as with magnesium and ferrous salts) 
from neutral solutions free from ammonium salts a part of the man-- 
ganese as the white hydroxide: 


MnCle+NH40OH @ Mn(OH)2+2NH4Cl. 


If sufficient ammonium chloride is present, ammonia causes no 
precipitation (cf. p. 94). The greater part of the manganese then 
remains in solution as manganous chloride, but a small amount exists as 
the hydroxide. On standing in the air, this dissolved hydroxide is 
changed slowly. into the more difficultly soluble manganous acid, which is 
deposited in brown flocks. The condition of equilibrium in the solution 
is thereby disturbed, and in order to restore it more hydroxide is formed, 
and the reaction continues in this way until finally all of the manganese 
may be precipitated. This fact must be considered in the separation 
of manganese from ferric iron, aluminium, ete. If a solution of ferric 
and manganous chlorides contains sufficient ammonium chloride, none 
of the manganese and all of the iron will be precipitated on the addition 
of ammonia, but, if the solution stands in contact with the air, little 
by little the manganese will be precipitated. In effecting the separa- 
tion, therefore, an excess of ammonium chloride should be present, the 
solution boiled to remove the air as much as possible from the solution, 
then a slight excess of ammonia should be added and the solution fil- 
tered immediately. The separation even then is not quantitative, but 
is satisfactory for qualitative analysis. 


Oxidizing agents in the presence of ammonia cause the precipitation of 
manganese as H,MnO;. Bromine is ordinarily used as the oxidizing agent, but 


MANGANESE 165 


a number of precautions are necessary to accomplish the complete precipitation 
of manganese by means of bromine and ammonia. 

If a neutral solution of manganous salt is treated with bromine, the precipita- 
tion of manganese as H:Mn0Q; is always incomplete: 


Mnt +4 Br.+3H.O y ax 4 4Ht++2Br-+H.Mn0O,. 


The precipitation of the manganese can be made complete, in accordance with 
the mass action principle (p. 13) if the hydrogen ions formed in the reaction 
are neutralized; sodium bromide in neutral solution will not reduce H.MnO,, 
but hydrobromic acid will do so. The solution may be neutralized by caustic 


alkali, alkali carbonate, alkali acetate (cf. p. 45), or ammonia. A solution of 


manganous salt may contain a considerable excess of acetic acid and yet the 
manganese will be completely precipitated by bromide in the presence of 
sodium acetate. 

Ammonia is not altogether satisfactory as a neutralizing agent in this case 
because it reacts with bromine as well as with hydrobromic acid. When 
bromine is added to ammonia solution, a vigorous reaction takes place and 
nitrogen is evolved (cf. p, 92); 


8NH,OH+3Br, =6NH,Br+8H,.0+N, t 


When bromine is added to a solution of ammonium chloride a very slow oxida-: 
tion results and nitrogen gas is evolved little by little: 


2NH,*+3Br, @ 8Ht+6Br +N; fT. 


If sodium acetate is added to this solution, the reaction is accelerated greatly as 
a result of diminishing the concentration of the hydrogen ions and, by heating 
the solution, which increases the speed of the reaction and causes the rapid 
expulsion of nitrogen gas, all of the ammonium salt can be decomposed by 
means of a very slight excess of bromine. 

The presence of ammonium salts prevents the precipitation of manganese by 
bromine and sodium acetate; when all the ammonium salt has been oxidized to 
nitrogen, the manganese can be precipitated as H,MnQ;. 

The conditions are more favorable for the precipitation of manganese when 
the solution contains hydroxyl ions, as in an ammoniacal solution. The addi- 
tion of bromine to such a solution usually results in the immediate precipitation 
of some of the manganese but, as a result of the action of bromine on ammonia 
or ammonium salt, the solution usually becomes acid and the precipitation of 
the manganese is then incomplete. To precipitate all of the manganese by 
means of ammonia and bromine, it is best to proceed as follows: 

Dilute the solution to about 200 cc. and neutralize, if necessary, with am- 
monia. Add a little bromine water, a slight excess of ammonia and stir to 
promote the formation of a precipitate. Heat the solution, add a little more 
bromine water and make slightly ammoniacal again. In case the amount of the 
manganese precipitate is perceptibly increased by this last treatment with 
bromine and ammonia, repeat the operation. Filter off the precipitated 
H:Mn0O;, concentrate the solution somewhat by evaporation and again treat 
with bromine water and ammonia. Sometimes H,MnO, is precipitated on the 
sides of the vessel during evaporation. The treatment with bromine and 


166 REACTIONS OF THE METALS 


ammonia should be continued until a filtrate is obtained which will not give 
any more precipitate with these reagents, * 


3. Alkali Carbonates precipitate white manganous carbonate, 
Mntt*+C03" > MnCOs, 


which after long boiling is changed by the oxygen of the air into less- 
soluble, hydrated manganese dioxide: : 


2MnC03+2H20+02=2CO2 | +2H2MnOs. 


4. Ammonium Carbonate precipitates even in the presence of 
ammonium salts the white carbonate (difference from magnesium). 
5. Barium Carbonate produces a precipitate only in hot solutions. 
6. Sodium Phosphate precipitates white, tertiary manganous 
phosphate, 
4HPO4.-+3M nits 2H2PO04-+Mnsz(POx4)2, 


soluble in mineral acids and in acetic acid: 
Mn3(PO.4)2+2H*t @ 3Mn*t+2HPO.-. 


If to the boiling solution of this precipitate in acid an excess of 
ammonia is added, manganous ammonium phosphate will be precipi- 
tated, as with magnesium (see p. 95): 


| HPO.4z-+NH40H @ NH4t+ PO4=+ H20. 
Mntt-+NH,t+ PO4=+7H20 als Mn(NH,4)PO,4 . 7H2O. 


The precipitate consists of pink scales and is practically insoluble in 
water. 

7. Lead Peroxide and Concentrated Nitric Acid. (Volhard’s 
reaction).—If a solution containing only traces of manganese is boiled 
with lead peroxide and concentrated nitric acid, then diluted with 
water and the residue allowed to settle, the supernatant liquid acquires 
a distinct violet-red color, owing to the formation of permanganic acid: 


2Mn*+t-+5PbO2+6H*t > 5Pbt*+2H20+2HMn0Og. 


This extremely delicate reaction does not take place in the presence 
of much hydrochloric acid or chlorides, because the permanganiec acid 
is thereby destroyed:. 


2HMn0O4+14HCl=8H20+ MnCle+5Cle T . 


* The above explanation is given at length because of its importance in quantita- 
tive analysis. The facts upon which the explanation is based have been carefully 
verified by quantitative experiments performed in the laboratory of the translator. 





MANGANESE 167 


8. Sodium Bismuthate added to a cold solution of a manganous 
salt in dilute nitric acid (about sp.gr. 1.13) causes the formation of 
permanganic acid. The reagent, which corresponds approximately to 
the symbol NaBiOs, is prepared by fusing bismuth oxide with sodium 
peroxide; it is insoluble in water and the excess of reagent may be 
filtered off through asbestos after applying the test. The reaction may 
be expressed by the equation: 


2Mnt*t+5NaBi03+16Ht > 5Nat+5Bitt++7H2,0+2HMn0,. 


The test is extremely delicate when nothing is present that will 
react with the permanganate formed. An insoluble carbonaceous 
‘residue, such as remains after the solution of cast iron in acid, must be 
filtered off before adding the reagent. If the solution is heated, the 
permanganic acid breaks down and hydrated manganese dioxide is 
precipitated. 

Bismuth dioxide, BiOz2, may be used instead of sodium bismuthate. 

9. Ammonium Persulfate. If a hot solution of a manganous 
salt in either dilute sulfuric or nitric acid is treated with ammonium 
persulfate, all the manganese is gradually oxidized to the quadrivalent 
condition and a precipitate of hydrated manganese dioxide is formed: 


Mn*t+820.7+3H20 — 2804-+4H*+ HeoMnOs. 


If, however, the solution contains a trace of silver nitrate as catalyzer, 
then the oxidation goes farther and permanganic acid is formed: 


2Mn*t-+5S820s-+8H20 — 16Ht+10S8S04-+2Mn0.-. 


This reaction is quantitative for small amounts of manganese and 
in the absence of anything that will react with the permanganic acid.* 

10. Ammonium Sulfide precipitates from manganese solutions 
flesh-colored, hydrated manganese sulfide: 


Mn**+S~— MnS. 


On boiling with a large excess of ammonium sulfide it is changed 
into less hydrated green manganese sulfide of the formula, 3MnS+ H20. 

The solubility product of MnS is relatively large (cf. p. 22) and to 
precipitate all the manganese as sulfide an excess of S~ ions is necessary. 
The precipitate dissolves readily in dilute acid as a result of the removal 
of the S~ ions to form non-ionized hydrogen sulfide (cf. p. 47). 

11. Potassium Cyanide.—On adding potassium cyanide to a solu- 
tion of a manganous salt, a brown precipitate appears which dissolves 





* Cf. M. MarsHat, Z. anal. Chem., 48 (1904), 418, 655. 


168 REACTIONS OF THE METALS 


in an excess of potassium cyanide, forming a brown solution. On 
standing, or by heating the solution, a voluminous green precipitate of 
K[Mn(CN)3] is formed which is soluble in strong potassium cyanide 
solution: 


Mn**+2CN- @ Mn(CN)2, Mn(CN)2+KCN = K[Mn(CN)3] 
K[Mn(CN)3]+3CN~ = K*+[Mn(CN)6¢]-~. 


To keep the manganese in solution in the form of Mn(CN)¢ ~ ions, 
it is necessary to use an excess of potassium cyanide. If the concen- 
tration of the cyanide is diminished by dilution, some green K{Mn(CN)s] 
is formed and if the dilute cyanide solution is boiled, a precipitate of 
Mn(OH): results: 


Mn(CN)6 ~+2H20 @ 2HCN+4CN +Mn(OH)>. 


The stability of these complex cyanides, therefore, is much less than that 
of the corresponding nickel compounds (p. 175). -This permits an 
interesting method of separating nickel from manganese. 

Nickel sulfide is much less soluble than manganous sulfide (cf. 
p. 22) so that it is possible to precipitate nickel as sulfide in the presence 
of acetic acid and sodium acetate; under these conditions no manganese 
sulfide is formed. If, on the other hand, ammonium sulfide is added to — 
a hot, dilute solution containing the complex cyanides of nickel and 
manganese, the nickel will remain in solution and the precipitation of 
the manganese as sulfide will be complete: 


Mn(CN)¢ +S <6CN +MnS. 


In the presence of a large excess of potassium cyanide, however, the 
Mn(CN)¢5 ~ anion is so stable that none of the manganese is precipi- 
tated in the cold by ammonium sulfide. 

12. Potassium Chlorate. By boiling a solution of manganous salt 
in concentrated nitric acid with an excess of potassium chlorate, all 
of the manganese is precipitated as MnQz. 


Mn(NO3)2+2KCl1O3 ~Mn0O2+2KNO03+2Cl0sz. 


REACTIONS IN THE DRY WAY 


The bead of borax, or salt of phosphorus, is amethyst. red after heat- 
ing in the oxidizing flame with small amounts of manganese, almost 
brown with larger amounts, and can then be mistaken for the nickel 
bead. Heated in the reducing flame, the manganese bead becomes 
colorless, while the nickel bead appears gray. 


to mee oe 


MANGANESE 169 


On fusing any manganous compound with caustic alkali or alkali 
carbonate (on platinum foil) in the air, or, better still, in the presence 
of an oxidizing agent (such as potassium nitrate, potassium chlorate, 
etc.), a green melt is obtained, owing to the formation of the alkali 
salt of manganic acid, as is shown by the following equations: 


Mn0O+Na2C03+02=CO2 | +NazMnO,, 

Mn0O2+ NuzC03+0 =CO2 | +Na2MnOa, 
Mn203+2Naz2C03+30 =2CO2 7 +2Na2MnOz, 
Mn304+3Na2CO03+50 =3CQO2 | +3Na2MnOz, 
MnS04+2Na2C03+ O02 =2COz2 | +NaeSO4+ NazMnO,. 


The oxygen comes either from the air or from the nitrate or chlorate: 
KNOz3=KNO2+0, KClO3=KCI+30. 


This reaction is exceedingly delicate; a fraction of a milligram 
of any manganese compound can be recognized by the formation of 


this green color. 
By ignition in the air the oxides of manganese are changed to Mn30a: 


3Mn0+0 = Mng0Quz, 3MnO2 = Mnsg04+O02 T ; 
6Mne203 =4Mn304+ O2 T ‘ 


B. Manganic and Permanganic Acids 


The free manganic acid has never been isolated. If we attempt 
to form it from the green melt of the alkali manganate by the addition 
of acid, permanganic acid and hydrated manganese dioxide will be 
obtained; a part of the unstable manganic acid oxidizes another part 
of the same to permanganic acid, while the oxidizing part is itself 
reduced to hydrated manganese dioxide: 


3HeMnO4 = 2HMn0O4+HeMnO03+ H20. 


This transformation takes place so readily that the green solution 
of the manganate is changed to a reddish-violet solution of a permangan- 
ate by simply standing in the air, with the help of the carbonic acid 
which the air always contains: 


3KeMn04+2C0e+ H20 = 2KeCO03+ H2Mn0O3 +2KMn0Osg. 


170 REACTIONS OF THE METALS 


; 


The reaction takes place much more rapidly, however, if a few 
drops of a strong acid are added. 


The oxidation of one molecule at the expense of another of the same 
kind is of quite common occurrence in chemistry. It always involves a 
loss in the available or free energy which the molecules originally possessed. 
The total energy possessed by any molecule can be considered to consist partly 
of free energy and partly of unavailable energy. A reaction that takes place 
spontaneously is always characterized by the fact that the free energy of the 
system is less afterward than it was before the reaction took place. The coi- 
dition with the smallest free energy is the most stable condition. 

It might be inferred that the most stable conditions are those having the 
smallest quantities of energy, but a little consideration shows that in promoting 
chemical reactions it is not so much the total energy as it is the available energy 
which comes into consideration. Thus, there is a vast amount of energy stored 
up in the heat of the ocean, but it is not available energy, because it is in sur- 
roundings at the same temperature. The air under ordinary atmospheric pres- 
sure could perform a great deal of work if it were brought in contact with a 
space in which a much lower gas-pressure prevailed, but otherwise the vast 
amount of energy is not available. 

In the changes that take place with any given element it is not necessarily 
true that an increase in the total energy will always involve an increase in the 
free energy associated with the element. The fact that the solution of a man- 
ganate, in which the valence of the manganese is six, decomposes readily 
indicates merely that the free energy in the system composed of quadrivalent 
and heptavalent manganese is less than the free energy involved in the system 
containing all the manganese in the hexavalent condition. Frequently the 
conditions are just the reverse and the most stable condition is one of inter- 
mediate valence. Thus permanganate and manganous salt react to form quad- 
rivalent manganese. A few reactions similar to the decomposition of HzMnO, 
will be given. 

Hypochlorites are changed, by warming the aqueous solution, into chlorate 
and chloride; one atom of chlorine is oxidized from the valence of one to the 
valence of five at the expense of two atoms of chlorine, which are reduced from 
a positive valence of one to a negative valence of one; 


3NaClO =NaClO;+2NaCl. 


Ignition of a chlorate causes the formation of a perchlorate, a chloride and 
free oxygen. Here one atom of chlorine is increased two in valence, oné 
atom of chlorine loses six charges and the remaining four charges cause the 
oxidation of two atoms of negatively charged oxygen: 


2NaClO; = NaClO,+NaCl+0, fT. 


Nitrous acid is changed in aqueous solution into nitric acid and nitric oxide, one 
atom of nitrogen gaining two charges and two similar atoms losing one 
charge: 


MANGANESE 171 


Hypophosphorous acid and also phosphorous acid can be changed into Boe: 
phoric acid and phosphine: 


2H;PO, =H,;P0O,+PH; f , 4H;PO, =3H;P0,+PH; T . 


Alkali thiosulfates and alkali sulfites are changed by ignition into sulfate 
and sulfide: 


4Na.8.0; = 3Na.SO.,+ NaS; ’ 4Na.SO; = 3Na.S80,+ NaS. 


Permanganic Acid, HMnO,, although much more stable than 
- manganic acid, is known only in aqueous solution; but the anhy- 
dride Mn207 has been isolated. On cautiously adding concentrated 
sulfuric acid to the cooled solution of a permanganate, oily drops of 
reddish-brown Mn207 separate out, which, however, on being warmed 
(the heat of reaction ‘is sufficient), explode with scintillation: 


2Mn207 +4H22SO4 =4MnS04+4H20+502 t . 


The salts of permanganic acid (the permanganates) are all soluble 
in water, with a reddish-violet color, and are very energetic oxidizing 
agents. In acid solution the heptavalent manganese in permanganate 
is usually reduced to bivalent manganous salt, but in alkaline, or nearly 
neutral, solutions manganese dioxide is the usual product. 

Oxidation in Acid Solution. Typical oxidation equations with 
permanganate have already been explained on p. 31. <A few of 
these reactions will be repeated here, but it will not be necessary to 
enter into further details concerning the method of balancing the equa- 
tions. To avoid exact repetition, the equations will be given with the 
entire molecules written instead of merely the ions involved, 


2KMn04+3H2804+ 10HCl = K2804-+2MnS804+8H20+5Cle J , 
2KMn04+3H2504+ 10HI = K2804-+-2MnS04+8H20 +52, 
2KMn04+3H2804+5He2S = Ke8O4+2MnS04+8H20-+5S8, 
8KMn04+12H2S04+5PH3 =4K2804+8MnS04+ 12H20+5H3PO4. 
2KMn04+6802+2H20 = 2KHS04+2MnS04+ Hed20¢. 


In this last equation the proportion of sulfate and of dithionic acid, 
H2S820¢, will vary with the temperature and concentration of the solu- 
tion: 
2K Mn04+5H2eC204+3H2S04 = Ko804+2MnS04+8H20+ 10C02 Lie 
2KMn04+5H202+4H2S04 = 2KHS01+2MnS01+8H20+502 7 , 
2KMn04+5KeC2064+ 14H2S04= 

= 12KHS04+2MnS0.i+8H20+ 10CO02+502 7. 


172 REACTIONS OF THE METALS 


Persulfuric acid, which is analogous to percarbonic acid, does not 
reduce a solution of a permanganate. An interesting reaction is 
that which takes place in nearly neutral solution between permanganate 
and manganous ions. The principal product is MnOz, which will carry 
down some of the bivalent manganese as manganous manganite unless 
an excess of some other ion is present which forms an insoluble man- 
ganite, e.g., zinc or calcium ions; 


2KMn0O4 +3MnSO4 + 2HeO = 2KHSO,4 -+- 5MnO> +H2S0.. 


Oxidation in Alkaline Solution—Many organic substances are 
oxidized by permanganates in alkaline solution with precipitation 
of manganese dioxide. Thus formic acid is oxidized to carbonic 
acid, ethyl alcohol to aldehyde and acetic acid, cellulose (paper) chiefly 
to oxalic acid, so that a. solution of a permanganate cannot be 
filtered through paper. By boiling a concentrated solution of potassium 
permanganate with concentrated potassium hydroxide, potassium 
manganate is formed with evolution of oxygen, and the color of the 
solution becomes green: 


4KMn04--4KOH = 4KoMn04-+2H20+0s> 1. 


By heating solid potassium permanganate to 240° C., potassium 
manganate is formed, also with evolution of oxygen: 


2KMn04= KeMn0O4+Mn0Oe +O2 T . 


NICKEL, Ni. At. Wt. 58.68 


Sp. Gr. =8.9. M. Pt. =1452°C. 


Occurrence.—In the native state nickel occurs only in meteor- 
ites. It is most frequently found in combination with sulfur, arsenic, 
and antimony in regular and hexagonal crystallizing minerals, of which 
the following are the most important: 

Isometric System: Chloanthite, NiAse; gersdorffite, NiAsS; ullman- 
. nite, NiSbS. Hezagonal System: Niccolite, NigAse; breithauptite, 
NigSbe2; millerite, NigSe. 

Nickel also occurs as regular crystals of bunsenite, NiO, isomor- 
phous with periclasite, MgO, and manganosite, MnO; as garnierite 
or noumeite, He(NiMg)SiO4-+aq, a mineral occurring in New Cale- 
donia, from which pure nickel can be prepared; and finally as anna- 
bergite, Nig(AsO4)2-8H2O, isomorphous with erythrite. 


NICKEL | 173 


Metallic nickel possesses a silver-white color and is difficultly 
soluble in hydrochloric and sulfuric acids, but readily soluble in nitric 
acid. It forms two oxides: green nickelous oxide, NiO, and brownish 

black nickelic oxide, NigOs. 
By dissolving either of these oxides in acids, salts of bivalent nickel 
are always obtained: 


Ni0+2HCI=H20+NiCh, 
Ni.03+6HC1=3H20+2NiCle+Cl ft , 
INi203-+4H2804 = 4H20 + 4NiSO2+02 1 . 


Nickelous oxide behaves as a basic anhydride, but nickelic oxide 
acts as a peroxide and forms no salts. 

The crystallized salts of nickel and their aqueous solutions are 
green, but in the anhydrous condition they are usually yellow. Most 
of the salts are soluble in water; the sulfide, carbonate, and phosphate 
are insoluble. 


REACTIONS IN THE WET WAY 


1. Potassium Hydroxide precipitates apple-green nickelous hydrox- 
ide, 
Nitt+20H~— Ni(OH)s, 


insoluble in excess of the precipitant, readily soluble in acids. 
2. Ammonia precipitates (in neutral solutions free from ammo- 
nium salts) a green basic salt, 


2NiSO4+2N H40H = (NH4)2804+NieSO,4- (OH)2, 


soluble (with a blue color) in excess of ammonia, forming complex 
nickel ammonia ions (cf. p. 25). 


Ni2SO.- (OH)2+12NH3 — 2Ni(NH3)6"*+20H~+S04°-. 


In the presence of sufficient ammonium salt, ammonia produces 
no precipitate, as with magnesium, ferrous and manganous salts; potas- 
sium and sodium hydroxides, however, precipitate the green hydroxide 
(difference from cobalt, see p. 179). 

The anhydrous chloride and sulfate readily absorb ammonia, 
forming anhydrous nickel ammonium salts: 


NiCle+6NH3 =[Ni(NHs3)6]Cl2; NiSO4+6NH3 =[Ni(NHs3)6|SOx. 


174 REACTIONS OF THE METALS 


The nickel ammonia cations are very stable in the presence of 
an excess of ammonia. In pure water they are in equilibrium with 
a small quantity of nickel ions of dissolved ammonia, and ammonium 
hydroxide: : 


Ni(NH3)6** — Ni**+6NH3; NH3+H20 — NH4OH. 


3. Potassium and Sodium Carbonates precipitate apple-green 
nickel carbonate: 
Ni**+CO37 > NiCOs. 


4. Ammonium Carbonate behaves similarly, but the precipi- 
tate which is formed is soluble in an excess of the precipitant, form- 
ing nickel ammonia carbonate. 

5. Sodium Hypochlorite precipitates in the presence of alkalies 
all of the nickel as brownish-black nickelic hydroxide, Ni(OH)s3. 
Nickelous hydroxide is first formed by the alkali present, but it is 
then oxidized by the hypochlorite to nickelic hydroxide: 


2Nit*+40H-+ClO~+H20 — 2Ni(OH)3+Cl-. 


On adding chlorine or bromine to the nickel solution to which alkali 
has been added, nickelic hydroxide is likewise formed: 


2Ni(OH)2+20H +Cle — 2Cl"-+2Ni(OH)s. 


6. Barium Carbonate produces in the cold no precipitation; but*by 
continued boiling, all of the nickel is thrown down as basic carbonate. 

7. Hydrogen Sulfide precipitates no nickel from solutions which 
contain mineral acid or much acetic acid; but from solutions slightly 
acid with acetic acid and containing an alkali acetate, all the nickel is 
precipitated as the black sulfide: 


Nit *+2C2H302-+H2S — 2HC2H302+NiS. 


.8. Ammonium Sulfide precipitates from neutral solutions the 
nickel as sulfide: 


NiCl2+(NH4)28 = 2NH4CI+NiS. 


Nickel sulfide has a marked tendency to form colloidal solutions 
of a dark-brown color, especially in the presence of ammonia or a con- 
siderable excess of ammonium sulfide. By making the brown solution 
slightly acid with acetic acid and boiling, the hydrosole is coagulated 
and can be removed by filtration. The presence of ammonium salts 
also favors the coagulation of the hydrosole, 


NICKEL 175 


If it is desired to precipitate the nickel as sulfide from an ammoniacal 
solution, it is best to make the solution very slightly acid, add a little ammonium 
chloride unless considerable is already present, heat to boiling, add colorless 
ammonium sulfide drop by drop until no further precipitation takes place, 
and then add 0.5 to 1 cc. of the reagent in excess. The nickel sulfide thus 
obtained can be filtered without difficulty and the filtrate is free from nickel. 
During the filtration care should be taken to keep the filter well filled with liquid 
to prevent the oxidation of the precipitate, which takes place readily on exposure 
to the air. To wash the precipitate it is well to use a hot, 5 to 10 per cent 
ammonium chloride solution to which a little colorless ammonium sulfide 
has been added. The washing can also be effected with hydrogen sulfide » 
water without there being any danger of hydrosol formation. 

Nickel sulfide is difficultly soluble in dilute mineral acids, readily soluble, 
however, in strong nitric acid or in aqua regia, with separation of sulfur: 


38NiS+6HCIl+2HNO; =3NiCl.+2NO T +4H,0+8S. 


The sulfur usually separates out as a black film. This is caused by the 
sulfur first melting, owing to the heat of reaction, enclosing small particles of 
the black sulfide and protecting them from the action of the acid. By continued 
action of the acid all the sulfide is dissolved, and the sulfur remains as yellow 
drops, which are oxidized little by little to sulfuric acid, 


S+2HNO, =H,.S0,+2NO tT e 


Nickel and cobalt sulfides, though not precipitated by hydrogen sulfide 
from a dilute, hydrochloric acid solution, dissolve with difficulty in a much 
stronger acid. This is perhaps due to the fact that these sulfides exist in two 
allotropic forms of different solubilities. The sulfide first precipitated is readily 
soluble in acid, but on standing it becomes changed into a much more insoluble 
condition. Most schemes of qualitative analysis are based upon this behavior; 
the nickel and cobalt are separated from zinc and manganese by treating the 
ammonium sulfide precipitate with cold, dilute hydrochloric acid. In such cases 
some nickel and cobalt always passes into solution and the quantity dissolved 
may be much larger than is ordinarily assumed. If the surface exposed to the 
action of the acid is large, or if left in a finely divided state by the dissolving out 
of other sulfide, a considerable quantity of nickel passes into solution in a com- 
paratively short time. The reverse reaction, the precipitation of nickel sul- 
fide by hydrogen sulfide in very dilute acid solution, also takes place very 
slowly but continuously.* 


10. Potassium Cyanide produces a light-green precipitate of 
nickelous cyanide readily soluble in an excess of the precipitant, forming 
potassium nickelocyanide: 


Nitt+2CN- — Ni(CN)2; Ni(CN)2+2CN~ — [Ni(CN)a]~. 
The Ni(CN)4~ anion is a stable complex, but it is decomposed by 





* Cf. Noyrs, Bray and Spear, J. Am. Chem. Soc., 1908. 


176 REACTIONS OF THE METALS. 


the addition of acid. This is because He[Ni(CN),] is a very weak acid 
and, like carbonic acid, is unstabie: | 


Ni(CN)47-+2H* = He[Ni(CN)4]; He[Ni(CN)4] @ 2HCN+Ni(CN)o. 


Hydrogen Sulfide in an Alkaline Solution containing a tartrate 
gives a clear brown solution. This is a very characteristic test for 
nickel and it enables the detection of as little as 0.2 mg. of nickel in 
20 cc. of solution; the exact condition of the nickel in this solution is 
not known. The main function of the tartrate is to form a complex 
anion with nickel and thus prevent the precipitation of nickel hydrox- 
ide. The characteristic brown solution is not obtained until the solu- 
tion is nearly saturated with hydrogen sulfide. | 

In alkaline solutions, containing an excess of cyanide, the [Ni(CN)4]~ 
anion is not dissociated to a sufficient extent into simple Nit* cations 
to give a precipitate with ammonium sulfide (difference from man- 
ganese and zinc) but it is readily decomposed by an oxidizing agent 
such as chlorine, bromine, hypochlorite or hypobromite 


QNi(CN)4]-+40H-+9BrO-+H20 — 2Ni(OH)3+8CNO~+9Br-, 
or 


QINi(CN)4]-+60H~+9Clz — 2Ni(OH)3+8CNCI+10CI-. 


In the above reactions, first of all any excess alkali cyanide present is Oxi- 
dized to cyanate or to CNCl. In the absence of an excess of cyanide ions, the 
[Ni(CN),]~ begins to dissociate, [Ni(CN),J- ~ N iY +4CN~ , and as fast as the 
ions are formed, they both become oxidized. 

When cobalt i ions are treated with an excess of cyanide ions a cobaltocyanide 
anion, [Co(CN) 4), is formed which is readily oxidized even by the oxygen 
of the air to form cobalticyanide ions, Co(CN),>~. These anions are so stable 
that they are not decomposed by treatment with oxidizing agents in alkaline 
solution so that no cobalt is precipitated with Ni(OH); by the above treatment. 
It is possible to detect 0. 2 mg. nickel in the presence of 300 mg. of cobalt by this 
reaction. 

The test can be applied to the solution obtained by dissolving the sulfides 
of nickel and cobalt in aqua regia. The test can be made as follows: Evaporate 
the solution nearly to dryness to expel most of the acid, add about 5 ee. of water — 
and then sodium hydroxide solution, drop by drop, until the solution is 
neutral or until a permanent precipitate is produced. Add potassium cyanide 
solution, a few drops at a time, until all or nearly all of any precipitated cyanide 
redissolves. Then add 0.5 to 3 ec. more of potassium cyanide solution accord- 
ing to the probable amount of nickel and cobalt present. Heat to 50° or 60° 
in an open dish for five minutes, or longer if the solution has not become light- 
colored. This serves to oxidize the cobaltocyanide to cobalticyanide. Filter 
off and reject any small precipitate that may remain. Add about 3 cc. of 2N 
sodium hydroxide solution and conduct chlorine into the solution in the cold or 
add a little bromine water. A precipitate of Ni(OH); should form within five 
minutes, 


NICKEL 177 


11. Sodium Phosphate precipitates apple-green nickel phosphate, 
3Nit +4 4HPO,= — 2HePO4, +Nisz (PO4) 2) 


readily soluble in acids, even acetic acid. 

12. Potassium Nitrite produces in dilute nickel solutions no pre- 
cipitate (difference from cobalt). In very concentrated solutions a 
brownish-red precipitate of Ni(NO)2-4KNOz is thrown down; in the 
presence of alkaline earth salts a yellow crystalline precipitate is formed; 
e.g., Ni(NO2)2:Ba(NO2)2:2KNOz2, which is very difficultly soluble 
in cold water, but readily soluble in boiling water, with a green color. 

18. Dimethylglyoxime. The reagent is prepared by dissolving 1 gm. 
of the solid in 100 cc. of 98 per cent alcohol. If a little of the 
reagent is added to a solution of a nickel salt, then ammonia to slightly 
alkaline reaction, and the solution is boiled, a red crystalline precipitate 
of the nickel salt of dimethylglyoxime is formed: 


CH3—C = NOH 
2 


CH;—C=NOH 


If the quantity of nickel present is very small, at first a yellowish 
solution is obtained from which, on cooling, red needles are deposited. 
According to L. Tschugaeff,* who first proposed this qualitative test, 
the presence of one part of nickel can be detected in the presence of 
400,000 parts of water. The reaction is not influenced by the presence 
of ten times as much cobalt; when a larger proportion of cobalt is 
present, the following procedure is followed, 


4-NiClo-+-2NH3=2NH,Cl+ (CsH4N,0.) Ni. 


Detection of Traces of Nickel in Cobalt Salts 


Add strong ammonia to the solution of the cobalt salt until a clear solution is 
obtained, then add a few cubic centimeters of hydrogen peroxide and boil the 
solution a few minutes to decompose the excess of this reagent. Then add the 
dimethylglyoxime and again bring the solution to a boil. A very small quan- 
tity of nickel causes a red scum to form and the glass sides of the beaker become 
coated with a film of red crystals. With smaller amounts of nickel the color 
is best observed upon the filter through which the solution is poured and the 
residue washed with hot water. 

The above reaction is the most sensitive test known for detecting nickel 
in the presence of cobalt. 

v. Fortinit recommends the following method for detecting nickel in alloys: 
Heat the metal, from which all greasy or oily matter has been removed, at one 
place with the oxidizing flame of the blowpipe. Cool and moisten with one 





* Ber. 38 (1905), 2520. 
t Chem. Zitg.(1912), 1461. 


178 REACTIONS OF THE METALS 


drop of a solution of 0.5 gm. dimethylglvoxime in 5 cc. aleohol and 5 ec. concen- 
trated ammonium hydroxide. A red spot at once appears if nickel is pres- 
ent. When copper is present, the nickel test is obtained before the copper 
ammonia color is visible. 


REACTIONS IN THE DRY WAY 
\ 


The borax, or sodium metaphosphate, bead is brown in the oxi- 
dizing flame, almost the same shade as the strongly saturated man- 
ganese bead; in the reducing flame the bead becomes gray, due to the 
formation_of some metallic nickel. On looking at the bead through 
the microscope the finely divided metal can be seen suspended in the 
colorless glass. | 


On heating nickel salts with sodium carbonate on charcoal a gray — 


scale of metallic nickel is obtained. This reaction is best performed 
with the charcoal stick, as described on p. 65. The magnetic metal 
obtained in this way is placed on a piece of filter-paper, dissolved 
in nitric acid, a drop of concentrated hydrochloric acid is added, and 
the paper carefully dried by moving it back and forth over the flame. 


If nickel is present the paper appears greenish (colorless with very small © 


amounts of nickel), or bluish if cobalt is also present. The paper is now 
moistened (where the nickel is) with caustic soda or potash, and is then 
held in bromine vapors, which are obtained by shaking some bromine 
water in a wide-mouthed flask. 

If nickel or cobalt is present, a black spot will be formed by the 
above treatment, consisting of the hydroxide of the trivalent metal 
(p. 174). The blackening often does not appear at first; in this case 
the paper is moistened once more with potassium hydroxide and again 
treated with bromine. The spot will now appear if nickel is present. 


COBALT, Co. At. Wt. 58.97 
Sp. Gr.=8.5. M. Pt. = 1490° 
Occurrence.—Like nickel, native cobalt is found only in meteorites. 


It occurs in the earth’s crust chiefly as sulfide, arsenide, and as salts 
of thioarsenious and thioantimonous acids; it is almost always accom- 


panied by nickel and iron. The most important ores are smaltite, — 


CoAsz, isometric; cobaltite, CoAsS, isometric; skutterudite, CoAss, 
isometric; erythrite, Cog(AsO4)2-8H20, monoclinic, isomorphous with 
vivianite, Fe3(PO4)2-8H20, and with annabergite, Nis(AsO4)2:8H20. 

Metallic cobalt is steel gray, dissolves much more readily in dilute 
acids than nickel, and is, like the latter, magnetic. Cobalt forms, like 
iron, three oxides: cobaltous oxide, CoO; cobaltous cobaltic oxide, 
Co304; cobaltic oxide, Co2Qz. 


ee is ee 


COBALT 179 


By dissolving these three oxides in acids, salts derived from cobalt- 
ous oxide are always obtained, containing bivalent cobalt: 


CoO +2HCl=H20 + CoCle, 
Co203 +6HCl = 3H2O+ 2CoCle +- Cle t ; 
Co304+8HCl=4H20+3CoCle+Cle fT . 


Simple cobaltic salts are unknown, but many complex compounds 
exist with trivalent cobalt, as, for example, potassium cobaltinitrite, 
potassium cobalticyanide, and numerous cobalti-ammonia deriva- 
tives. 

- Cobaltous compounds in a crystallized state (as well as in aqueous 
solution) are pink, in the anhydrous condition yellow or green, and 
blue in aqueous solutions in the presence of hydrochloric acid. The 
solubility reactions of cobaltous salts are similar to those of manganese 
and nickel. 

| REACTIONS IN THE WET WAY 


1. Potassium or Sodium Hydroxide precipitates in the cold a blue 


basic salt: | 
Cott+Cl-+0OH- — Co(OH)CI, 


which on warming is further decomposed by hydroxyl ions forming 
pink cobaltous hydroxide: : 


Co(OH)C1+OH- — Cl-+Co(OH)>. 


In the case of a moderately concentrated solution of the alkali the 
precipitate of pink cobaltous hydroxide is often produced in the cold, 
sometimes only after standing some time. The rapidity of the reaction 
depends entirely upon the concentration of the alkali. 

Cobaltous hydroxide gradually turns brown in contact with the air, 
going over into cobaltic hydroxide: 


2Co(OH)2+H20+0 — 2Co(OH)s. 


In this respect cobalt behaves similarly to iron and manganese, 
and differs from nickel, for the hydroxide of the latter is not oxidized 
by atmospheric oxygen. 

On adding chlorine, bromine, hypochlorites, hydrogen peroxide, 
etc., to an alkaline solution containing cobaltous hydroxide, cobal- 
tic hydroxide is immediately formed, as with nickel and manganese, 


2Co(OH)2+20H-+Cle = 2CI-+2Co(OH)s, 
2CO(OH)2-+HOH+0Cl-=Cl-+2Co(OH)s. 


180 REACTIONS OF THE METALS 


From ammoniacal cobalt solutions the above oxidizing agents 
cause no precipitation, but merely a red coloration; the addition of 
potassium hydroxide then causes no precipitation (difference from 
nickel). 


Cobaltous hydroxide, Co(OH).2, behaves under some conditions as a weak 
acid, for on adding to a cobaltous solution a very concentrated solution of KOH 
or NaOH the precipitate at first produced dissolves with a blue color* similar 
to copper. By the addition of Rochelle salts to this blue cobalt solution the 
color either disappears almost entirely or becomes a pale pink, while the similarly 
treated copper solution becomes more intensely blue. By the addition of 
potassium cyanide to the blue cobalt solution it becomes yellow, and in contact 
with air turns intensely brown. A copper solution would be decolorized by the 
addition of potassium cyanide. 

By pouring a little cobalt solution (or adding a little solid cobalt carbonate) 
into a concentrated solution of caustic soda or potash, to which a little glycerol 
has been added, a blue solution is formed (the color being intensified by warm- 
ing), which after standing some time in the air, or immediately upon the addition 
of hydrogen peroxide, becomes a beautiful green. 


2. Ammonia precipitates, in the absence of ammonium salts, a 
blue basic salt, soluble, however, in excess of ammonium chloride. 
Ammonia, therefore, produces no precipitate in solutions which contain 
sufficient ammonium chloride. The dirty-yellow, ammoniacal solution 
is little by little turned reddish on exposure to the air, owing to the 
formation of very stable cobalti-ammonia derivatives. 

3. Alkali Carbonates produce a reddish precipitate of basic salt 
of varying composition. 

4. Ammonium Carbonate also precipitates a reddish basic salt, 
soluble, however, in excess. 

5. Barium Carbonate precipitates no cobalt in the cold and out 
of contact with air, but on exposure to the air cobaltic hydroxide is 
gradually thrown down. The precipitation takes place much more 
quickly on the addition of hypocblorites or hydrogen peroxide: 


2Cot ++2BaCO3+3H20+0CI — 2Ba*t+Cl-+ 2Co(OH)3+2CO2 fF 


_ If the solution is heated to boiling, all of the cobalt is precipitated 
as a basic salt, even out of contact with the air. 

6. Hydrogen Sulfide produces no precipitate in solutions con- 
taining mineral acids. In neutral solutions containing an alkali - 
acetate all of the cobalt is precipitated as black sulfide. 

7. Ammonium Sulfide precipitates black cobalt sulfide, 


Cott+S8= — CoS, 
* Ep. Donatn, Z. anal. Chem., 40 (1901), 137. 


COBALT 181 


insoluble in ammonium sulfide, acetic acid, and very dilute hydro- 
chloric acid (cf. p. 175); soluble in concentrated nitric acid and aqua 
regia, with separation of sulfur: 


3CoS+8HNO3 =4H20+2NO 7 +3S+3Co(NOs)>o. 


By continued action of strong nitric acid all the sulfur goes into 
soiution as sulfuric acid. 


The addition of an oxidizing agent always helps an acid to dissolve an insolu- 
ble sulfide. The solution in contact with a sulfide precipitate at first contains 
enough sulfide ions to satisfy the solubility product of the sulfide. When 
hydrogen ions are added nonionized hydrogen sulfide is formed unless the 
solubility-product of the sulfide is so small that less sulfide ions are present 
than would be formed by the ionization of H.S. If an appreciable quantity 
of HS is formed, it can be expelled as a gas and the sulfide will dissolve. Some- 
times, however, this takes place very slowly and then the addition of an oxidiz- 
ing agent is necessary. ‘The sulfide ions in solution are oxidized to free sulfur. 
The solubility-product of the sulfide is no longer reached in solution, for as fast 
as a little of the substance dissolves the sulfide ions are oxidized. 


8. Potassium Cyanide produces in neutral solutions a reddish- 
brown precipitate, soluble in excess of potassium cyanide in the cold, 
with a brown color, forming potassium cobaltocyanide: 


Cott+2CN~ — Co(CN)2; CoCN2+4CN- — [Co(CN)6¢]"== 


On warming the brown solution for some time it becomes bright 
yellow and reacts alkaline; it now contains potassium cobalticyanide, 
of analogous composition to potassium ferricyanide. The formation 
of the cobaltic salt takes place with the help of atmospheric oxygen: 


2[Co(CN).6]~>~+O+H20 — 2[Co(CN).6]>+20H . 


The reaction takes place more quickly in the presence of chlorine, 
bromine, hypochlorites, etc.: 


2[Co(CN)¢]*=+Cle — 2[Co(CN)¢]=+2CI-. 


An excess of chlorine, bromine, etc., does not decompose the cobaltic 
salt (difference from nickel). 

The cobalticyanide anion is much more stable than the cobalto- 
- eyanide anion. By adding hydrochloric acid to the brown solution of 
potassium cobaltocyanide, hydrogen cyanide will be set free and yellow 
cobaltous cyanide formed, 


[Co(CN).]""+4Ht — 4HCN-+Co(CN)o, 


while potassium cobalticyanide is not decomposed by hydrochloric acid. 


182 REACTIONS OF THE METALS 


Potassium cobalticyanide forms, with most of the heavy metals, 
difficultly soluble or insoluble salts possessing characteristic colors. 
Thus, it produces with cobaltous salts pink cobaltous cobalticyanide: 


2[Co(CN)6]=+3Co** — Cos[Co(CN).le, 


and with nickel salts greenish nickelous cobalticyanide. 
If, therefore, a cobalt solution contains nickel it gives, on precipi- — 

tating and redissolving with potassium cyanide, boiling, and adding 

hydrochloric acid, a greenish precipitate of nickelous cobalticyanide: 


2[Co(CN)6}" 4-3[Ni(CN)4] +12Ht — 12HCN+Nis3[Co(CN)glo. 


9. Potassium Nitrite produces in concentrated solutions of cobalt 
salts, with the addition of acetic acid, an immediate precipitation of 
yellow, crystalline potassium cobaltinitrite. If the solution is dilute, 
the precipitate appears only after standing for some time, but more 
quickly on rubbing the sides of the beaker. 


Cot t-+-7NO. +3Kt+2Ht — NO+H20 +Ks[Co(NO2) 6]. 


This reaction offers an excellent means for detecting the presence 
of cobalt in nickel salts. 

10. Ammonium Thiocyanate (Vogel’s reaction).* If a concen- 
trated solution of ammonium thiocyanate is added to a cobaltous solu- 
tion, the latter becomes a beautiful blue, owing to the formation of 
ammonium cobaltothiocyanate: 


Cott+4CNS~+2NH,4t > (NH4)2/Co(CNS)z]. 


On adding water the blue color disappears and the pink color of the 
cobaltous salt takes its place. If amyl alcohol is added (or a mixture 
of equal parts amyl alcohol and ether), and the solution shaken, the 
upper alcoholic layer is colored blue. This reaction is so sensitive 
that the blue color is recognizable when the solution contains only 74> 
of a milligram of cobalt. The blue solution also shows a characteristic 
absorption spectrum.{ Nickel salts produce no coloration of the amyl 
alcohol. If, however, iron is present, red Fe(CNS)3 is formed, which 
likewise colors the amyl alcohol, making the blue color (due to the 





* Ber., 12, 2314; Treapwett, Z. anorg. Chem., 26 (1901), 105. 

+ T. T. MorreEtt first showed that cobalt salts give a blue color with ammonium 
thiocyanate, disappearing on the addition of water, but reappearing when alcohol 
is added. Z. anal. Chem., 16, 251. 

t Wotrr, Z. anal. Chem: 18, 58. 


COBALT 183 


cobalt) very indistinct, so that, under some conditions, it can no longer 
be detected. If, however, 2 or 3 cc. of concentrated ammonium acetate 
solution and 2 or 3 drops of 50 per cent. tartaric acid solution are added, 
the red color produced by Fe(CNS)3 will disappear and the blue color 
of the cobalt compound will be seen. 


The blue color is probably that of undissociated (NH,)s[Co(CNS),]. When 
the solution is diluted, the salt is ionized and the complex anion also is in 
equilibrium with cobalt ions, the more dilute the solution, the greater the 
ionization (cf. p. 19). The alcohol and ether probably dissolve only the un- 
dissociated (NH,).[Co(CNS),] and this is evidently present to some extent in 
the aqueous solution, although its color is obscured by that of cobaltous ions. 


11. Ether Saturated with Hydrogen Chloride does not precipitate 
an anhydrous cobaltous salt, as in the case of nickel, but will dissolve 
‘the blue, anhydrous cobaltous chloride. This furnishes a basis of a 
method for separating nickel and cobalt. 

12. a-Nitroso-6-naphthol, CjoHe6(NO)OH, produces a voluminous, 
purple red precipitate of cobalti-nitroso-8-naphtol, [Ci9H¢(NO)O]3Co, 
which is insoluble in cold, dilute nitric or hydrochloric acid.* 


This reagent serves not only for qualitative purposes, but can also be used 
for the quantitative determination of cobalt in the presence of nickel: The 
test may be applied conveniently to the solution obtained in the usual qualita- 
tive scheme after the removal of all metals except nickel and cobalt.. A part 
of the solution may be used for the sensitive nickel test with dimethylglyoxime 
(p. 177) and the remainder used for the cobalt test. 

Dilute the solution to about 50 cc., add 4 cc. of 6N hydrochloric acid and 
20 ec. of 6N acetic acid. Heat, add 50.cc. of a saturated solution of nitroso- 
6-naphthol and boil in 50 per cent acetic acid. If as much as 0.1 mg. of cobalt 
is present, a red precipitate or turbidity is obtained even in the presence of 250 
mg. of nickel. When more than 150 mg, of nickel is present, however, some of 
the brownish-yellow nickel compound, (CioHs(NO)O).Ni, will precipitate after 
the solution cools. 

The reagent used in this test should be freshly prepared. Nitroso- 
p-naphthol gradually decomposes on standing in the air and changes from yellow 
to brown or even black in color. It can be purified by dissolving in hot sodium 
carbonate, filtering and reprecipitating with sulfuric acid. For all ordinary 
purposes the saturated solution in 50 per cent acetic acid is most suitable. 
The cobalt test can be made more delicate by adding an equal volume of alcohol 
to the test and, for detecting traces of cobalt, an aqueous solution of the organic 
substance can be used, but as 5000 cc. of water are required to dissolve 1 gm. of 
the nitroso-s-naphthol, it is evident that the aqueous solution is not suitable 
when much cobalt is present. An excess of the reagent is required, as a part of 
it is used to oxidize the cobalt to the trivalent condition. 

Copper gives a characteristic coffee-brown precipitate with the reagent 
and it is possible to separate copper from leagl, cadmium, etc., by means of it. 





*TLInskI and v. Knorre, Ber., 18, 699 (1885), 


184 REACTIONS OF THE METALS 


Ferric iron gives a brownish-black precipitate which is insoluble enough to 
serve as a means of separating iron from aluminium, manganese, etc. Ferrous 
iron also gives a greenish precipitate in neutral solutions. Of all these precipi- 
tates, however, the cobalt compound is the most characteristic and the least 
influenced by the presence of acid. Thus with the acidity recommended above, 
the presence of a little ferric or ferrous iron causes no disturbance. 


Detection of Traces of Cobalt in Nickel Salts 


To test a nickel salt for ‘cobalt, add a concentrated solution of 
ammonium thiocyanate to the solution of a considerable amount of 
the salt, a few cubic centimeters of a mixture of amyl alcohol and ether 
and shake the mixture. After the latter has been allowed to settle, 
if the upper alcohol-ether layer is colorless, then the nickel salt contains 
neither iron nor cobalt; if the layer is reddish, iron is present. In the 
latter case add 2 or 3 cc. of concentrated ammonium acetate solution 
and 2 or 3 drops of 50 per cent tartaric acid solution and shake again; 
if cobalt is present the alcohol-ether layer is now distinctly blue. 

Sometimes when very little cobalt and considerable nickel is present it is 
hard to tell whether the amyl alcohol is colored blue or not. In such a case 
pour the solution into a separatory funnel and draw off the lower layer contain- 
ing the green nickelous solution. Add a little more ammonium thiocyanate 
solution to the amyl alcohol, 1 cc. of ammonium acetate solution, 1 drop of 


tartaric acid solution and shake again. The blue color should now appear if 
any cobalt is present. 


REACTIONS IN THE DRY WAY 


The bead produced by borax or sodium metaphosphate is blue in 
both the oxidizing and reducing flames. By holding the bead in the 
upper reducing flame for a long time it is possible to reduce the cobalt 
to metal, when it appears, like nickel, gray. 

On the charcoal stick cobalt compounds yield gray metallic cobalt, 
which can be removed by means of a magnetized knife-blade, as 
described on p. 66, placed on filter-paper, dissolved in hydrochloric 
acid and dried. ‘The paper is then colored blue by cobalt (difference 
from nickel). If, now, sodium hydroxide is added and the paper 
exposed to the action of bromine vapors, black cobaltic hydroxide, 
Co(OH)s3, is formed. 


ZINC, Zn. At. Wt. 65.37 


Sp. Gr.=6.9. M. Pt.=419°. B. Pt.=916°. 


Occurrence.—Smithsonite, ZnCOg, isomorphous with calcite, 
CaCOs, etc.; sphalerite, ZnS, isometric; calamine, Zn2Si04+He20, 
orthorhombic, hemimorphic; zincite, ZnO, hexagonal; and franklinite, 
(FeO2)2(Fe,Mn,Zn), isometric. : 


ZINC 185 


The most important zine ore is sphalerite, ZnS. Sulfide of zine 
is dimorphic and is also found as wurtzite, which crystallizes in the 
hexagonal system. 

Metallic zine is bluish white. At low temperatures and at about 
200° C. it is so brittle that it can be pulverized, but at 110°-150° C. 
it is ductile and can be drawn out into wire and rolled into foil. 

Zinc, as its position in the electromotive series would indicate (p. 41) 
dissolves readily in all acids; in hydrochloric, sulfuric, and acetic 
acids with evolution of hydrogen: 


Zn+2H* > Zn**+Hp f. 


Zinc is such a strong reducing agent that it easily reduces nitric acid, 
the extent of the reduction depending upon the concentration of the 
acid. With very concentrated acid, some NOz is obtained, while dilute 
acid is reduced to ammonium nitrate. In concentrated acid the prin- 
' cipal product is nitric oxide, NO. 


8Zn+8HNO3 — 3Zn* *+6NO3° +4H20+2NO f¢. 
4Zn+10HNO3 — 4Zn* *+8N03°+NH4NO3+3H20. 


- Like aluminium, zine dissolves in caustic soda or potash, with 
evolution of hydrogen and the formation of a zincate: 


Zn+20H~— ZnOo-+He 1T. 


Zine forms only one oxide, ZnO. It is a white infusible powder, 
which becomes yellow when heated, but turns white again on cooling. 
Zinc oxide dissolves readily in acids, forming zinc salts: 


ZnO +H22SO4 = H20 +7ZnSOx. 


There exists only one series of zinc salts, and the zinc is always 
bivalent. Most of the salts are white. The chloride, nitrate, sul- 
fate, and acetate are soluble in water; the remainder dissolve readily 
in mineral acids. 

REACTIONS IN THE WET WAY 


1. Potassium or Sodium Hydroxide precipitates white, gelatinous 
zinc hydroxide, easily soluble in excess of the precipitant, forming a 
— gzineate:* 


Zn+*++20OH- — Zn(OH)2; Zn(OH)2+OH- @ HZnO2 +H20. 





* According to Hantzscu the zinc is not present as zincate, but probably in 
colloidal solution. Z. anorg. Chem., 30, 289 (1902). In fairly concentrated solu- 
tions, however, it is certain that the zinc is present as zincate, for F. Forrster 
and O. Ginruer, Z. Electrochem., 6, 301 (1900), have isolated the compound, 
NaHZnO -3H,0, as needles with silky luster. 


186 REACTIONS OF THE METALS 


Zinc hydroxide, therefore, behaves sometimes as a base and some- 
times as an acid, like aluminium hydroxide. 

On boiling a diluted solution of a zincate, hydrolysis takes nlids 
and zine hydroxide is precipitated, but if the solution contains an excess 
of OH ions, there will be no precipitation. 

2. Ammonia precipitates from neutral solutions, free ‘from am- 
monium salts, zinc hydroxide, readily soluble in ammonium galts, 
as in the case of magnesium, nickel, manganese, or iron: 


Zn* ++9NHsOH ra 4 Zn(OH)>2 +2NHa,". 


Zinc hydroxide is also soluble in an excess of ammonia, due to the 
formation of complex zinc ammonia ions: 


Zn(OH)2+6NH3 — [Zn(NH3)6]* t+20H-. 


3. Alkali Carbonates precipitate a white, basic carbonate, of 
variable composition, as is the case with magnesium. 

4. Ammonium Carbonate does the same, except that the pre- 
cipitate is soluble in an excess of the reagent. ‘The presence of am- 
monium salts or of ammonia prevents the precipitation. 

5. Barium Carbonate precipitates no zinc in the cold, but on boiling 
all the zinc is precipitated as basic carbonate. 

6. Sodium Phosphate precipitates gelatinous, tertiary zine phos- 
phate, which soon becomes crystalline, and is soluble in ammonia and 
in acids: 

Zn* *++4HPO« — ZHePO4- +Zn3(PO4)o. 


In the presence of ammonium salt, the less soluble zinc ammonium 
phosphate is precipitated: 


Zn* +4+4.NH,4t+2HPO," — HePO, +ZnNHy4POs,. 


Both zine phosphate and zinc ammonium phosphate dissolve readily 
in dilute acids, owing to the formation of very slightly ionized HPO4’, 
and in dilute ammonia, owing to the formation of zinc ammonia cations. ° 
Acids, therefore, deprive the solution of PO4> anions by forming HPO’, 
and ammonia deprives the solution of zincions by forming [Zn(NH3)¢6]** 

7. Hydrogen Sulfide precipitates the zine as sulfide, from neutral 
solutions of a zinc salt: 


Zn+++H.S = ZnS+2H". 


The solubility product of zinc sulfide (p. 22) is about 1.2X107%%. At 25° 
the concentration of a saturated solution of hydrogen sulfide is about 0.1 molar 


ZINC 187 


and the ionization constant for the complete ionization, H.S = 2H-+5S, has been 
estimated to be 1.110~*8. The concentration of the sulfide ion in such a solu- 
tion is approximately 1.2X10~ molar equivalents per liter. The solubility- 
product of zine sulfide is evidently exceeded when the aqueous solution of 
a zinc salt is saturated with hydrogen sulfide and zine sulfide is precipitated. 
The mass-action principle applied to the complete ionization of hydrogen 
sulfide shows that the concentration of sulfur ions is inversely proportional 
- to the square of the concentration of the hydrogen ions. If the concentration 
of the hydrogen ions is increased one thousandfold, and this is approximately the 
case when the solution is tenth-normal with a mineral acid, the concentration 
of the sulfide ions from hydrogen sulfide is reduced to one-millionth of its value 
in pure water. 

The separation of the second group of metals from the third group is usually 
accomplished by passing hydrogen sulfide into a solution which is about 0.3- 
normal with hydrochloric or nitric acid. The concentration of the sulfide ion 
when such a solution is saturated with hydrogen sulfide at 25° is about 
1.1X10-*. To reach the solubility product of zine sulfide in 0.3-normal acid, 
the zinc ions should reach the concentration of about 0.11 mole per liter, or 
about 0.7 gm. per 100 ce. . 

If the zinc salt has a greater concentration than this, some zine sulfide 
should be precipitated by hydrogen sulfide in 0.3-normal acid solution. The 
precipitation would evidently be incomplete and, as more hydrogen ions are 
formed in solution from the hydrogen sulfide as a result of the sulfide precipita- 
tion, the ionization of the hydrogen sulfide continually tends to become less. 

If, however, considerable sodium or ammonium acetate is added to the acid 
solution, the concentration of the hydrogen ion becomes much smaller and the 
ionization of the hydrogen sulfide takes place to a greater extent. It is then 
possible to precipitate the zinc as sulfide so completely that less than one i mg. 
of zinc will remain in solution. 

Zinc sulfide dissolves readily in normal hydrochloric acid. The sulfur ions 
from the zinc sulfide enter into equilibrium with the hydrogen ions of the acid 
to form hydrogen sulfide. In normal acid solution, the concentration of sulfur 
ions from saturated hydrogen sulfide is about 1.2X10~%* and of sulfur ions 
from a saturated solution of ZnS in water about 3.5107". 


8. Ammonium Sulfide precipitates from neutral or alkaline solutions 
all the zine as amorphous sulfide: 


Zntt++S"- —> ZnS. 


Zinc sulfide is a precipitate hard to filter; it runs through the 
filter-paper, particularly on washing. This is a peculiarity of almost 
all metallic sulfides and of many other amorphous bodies, such as | 
aluminium hydroxide, titanic acid, tungstic acid, and many others. 
This is due to its tendency to form colloidal solutions (p. 58). The 
colloid can be precipitated by adding a concentrated salt solution or 
by boiling. 





* This value is merely an approximation, being derived by a rough calculation 
from values which are not very reliable. 


188 REACTIONS OF THE METALS 


In order, then, to obtain zine sulfide in a form which can be 
filtered, it is best precipitated from a boiling solution containing acetic 
acid and a considerable quantity of ammonium salts. The precipitate 
may be washed with a solution of ammonium chloride to which a little 
ammonium sulfide has been added. 

9. Potassium Cyanide produces a white precipitate of zine cyanide, 
soluble in an excess of the precipitant. 


Znt*++2CN~ @ Zn(CN)2; Zn(CN)2+2CN7~ — [Zn(CN)a]*. 
The zine-cyanide anion is decomposed by acids and by alkali sulfide: 
[Zn(CN)4]-+2Ht — Zn(CN)2+2HCN; 
[Zn(CN)4]- +S" — ZnS+4CN-. 


10. Potassium Ferrocyanide precipitates white zinc ferrocyanide, 
which is changed by-an excess of the potassium ferrocyanide into less 
soluble zinc-potassium ferrocyanide: 


¢ [Fe(CN) 6]=~+2Zn* + — Zno[Fe(CN)6] 
8Zn2[Fe(CN)«]+Ks[Fe(CN)¢] — 2K2Zn3[Fe(CN)6]2 


REACTIONS IN THE DRY WAY 


Heated with sodium carbonate on charcoal before the blowpipe, 
it is not possible to obtain metallic zine on account of its volatility; 
but an incrustation of oxide is obtained which is yellow while hot 
and white when cold. 

Zine oxide (or such compounds of zine as are changed over to oxide 
on ignition), when moistened with cobalt nitrate yields a green infusible 
mass—Rinnmann’s green. This reaction is performed exactly as with 
aluminium (p. 131). 


Separation of the Metals of Group III from the Alkalies and 
Alkaline Earths 


The separation of the members of the ammonium-sulfide group 
from the alkalies and alkaline earths is effected by means of ammo- 
nium sulfide and ammonium chloride. If, however, the solution 
contains phosphoric acid, oxalic acid, or considerable boric acid, the 
neutralization of the solutions will cause the precipitation of calcium, 
strontium, barium, and magnesium as phosphate, oxalate, or borate. 
The procedure to be followed when such acids are present will be given 
in Part IV of this book after the characteristic properties of the igi 
have been described in Part III. 


Taste V.—ANAtysis or Group III 1n ABSENCE OF PHOSPHATE. 


ANALYSIS OF GROUP III 


189 


Metuop A 





Solution may contain: Fett+, Fet+,UO,++, Altt+, Crt++, Mnt+, Znt+, 
Cot*+, Nit*+, also Groups IV and V. Add NH,OH and (NH;,)S. Filter and 

















examine filtrate for Groups IV and V. Treat ppt. with 2-normal HCl. (2) 
Residue: CoS, | Solution: Fet+, UO.++, Al+++, Crt++, Mn++, Zn++ [and some 
NiS. Test for | Ni]. Add NaOH in excess. (4) 
Ni and Co by 
bead test. Test | precipitate: Fe(OH);, Cr(OH)s, | Filtrate: AlO.-, HZnO.-. Add 
for Ni with) Na,U.07, Mn(OH)2 [and some| HCl and then NH,OH. (10) 
Nie de an Ni(OH).]. Dissolve in HCl, 
find for cobalt add NH,Cland NH,OH. (5) 
eae Roar Precipitate:|Filtrate: |Preci pi -| Filtrate: Zntt. 
anate. (3) Fe(OH);, Cr(OH);,}, Mn++] tate: AI(OH);| Acidify with 
(NH4)2U207. Dis-| (and Confirm by| HC.H;02. add 
solve in HCl| Nitt)|] Thénard’s| HS and con- 
and add excess| Test for| bluetest. (10)| firm Zn _ by 
(NH4)2CO3. (6) Mn by Rinnmann’s 
(NH4)2S green test. 
Precipitate:|Filtrate: et de (11) 
Fe(OH )s, | (NHs)U02(COs)s} Gp py fun 
Cr(OH)s.| Acidify with) bith 
Test for| HCl and test NaCO 
Fe with| with SS. 
K.iFe (CN )e. K4Fe(CN )<¢. KNO 
TestforCrby | (8) (9) ; 
fusion with 
Na,CO; and 
KNOs, ete. (7) 




















A number of excellent schemes have been devised for the analysis of this 
group and each has something in its favor. In this book it has seemed best 
not to attempt to decide upon any one scheme, but rather to treat the subject 
in a broader manner, partly because of the instructive value of studying several 
schemes and partly because one scheme is best under certain conditions and 
another scheme under different conditions. Thus it is a quite common practice 
to divide the whole ‘group into two minor groups, one containing ferric iron, 
aluminium, chromium, and uranium and the other containing manganese, 
nickel, cobalt and zinc. Such a scheme often works very nicely and enables one 
to arrive at proper conclusions quickly but, unfortunately, chromium when 
precipitated in this scheme has a marked tendency to carry down zine and 
magnesium with it, and this may result in the failure to detect zinc or magnesium 
during the subsequent examination. Chromium, however, on account of the 
color of its compounds, invariably betrays its presence before the actual test 
for chromium is made, and it is a very easy matter to modify the method some- 
what when chromium is present in order that zinc and magnesium will not be 
missed. It is unnecessary to use such a modified method when chromium is 
known to be absent. In most schemes of analysis the detection of uranium and 
titanium is not provided for in the analysis of this group. Titanium, although 


190 REACTIONS OF THE METALS 


present in most rocks, is usually found in very small quantities. When much 
titanium is present, this fact is known by the difficulty involved in getting the 
substance in solution and the tendency of the dilute acid solutions to hydrolyze 
and form precipitates of metatitanic acid, which are difficult to filter. When 
much titanium is present it is precipitated according to p. 158, §4. To detect 
small quantities of titanium, the hydrogen peroxide test (p. 159) is most suit- 
able. Uranium is of relatively rare occurrence, but its ores have become 
important since the discovery of radium in them. The detection of uranium 
will be included in the first of the schemes of analysis that follow. 


PROCEDURE 


1. Heat the neutral solution to boiling, add 5 ec. of normal ammonium 
chloride solution, if this salt is not already present, and ammonium sulfide 
solution drop by drop until no further precipitation takes place. Avoid add- 
ing an excess of ammonium sulfide on account of the danger of getting a turbid 
precipitate when nickel is present. (To avoid this danger it is well to pass 
hydrogen sulfide into the slightly ammoniacal solution instead of adding ammo- 
nium sulfide. The reason why nickel sulfide runs through the filter is partly . 
because ammonium polysulfide is present to some extent in the ammonium sul- 
fide reagent that is not freshly prepared.) Filter off the precipitated sulfides 
and wash promptly with hot water. If the moist sulfides are allowed to stand 
exposed to the air, some sulfate is formed by oxidation and this will dissolve in 
the wash water. Reject all but the first washings and use the filtrate for the 
analysis of the alkaline earth and alkali groups (p. 111). 

2. Digest the precipitated sulfides in a porcelain dish with cold, 2-normal 
hydrochloric acid, stirring until no more hydrogen sulfide is evolved. Filter 
off the residue, which consists chiefly of cobalt and nickel sulfides and wash with 
a little hydrochloric acid. Usually a partial oxidation of the sulfide takes place 
during this treatment with dilute acid and sulfur is formed which is likely to 
enclose a little sulfide that should dissolve in the acid. The fact that a residue 
remains is not, therefore positive proof of the presence of nickel or cobalt. 
Examine the solution by § 4. 

3. Test the residue for cobalt by heating a little of it in a borax bead; a 
blue bead shows cobalt. If the borax bead is brown, further test for nickel is 
unnecessary. If a blue bead was obtained, test for nickel by dissolving the 
precipitate in aqua regia, evaporating just to dryness, adding a little water 
and testing with dimethylglyoxime (p. 176). If a brown bead was obtained, 
dissolve the precipitate as just described and, to the solution freed from mineral 
acid, add a concentrated solution of potassium nitrite, acidify with acetic acid 
and allow the solution to stand at least ten minutes. A fine yellow precipitate 
of K;[Co(NO:).] shows that cobalt is present. 

4, Evaporate the solution obtained in (2) to a small volume, oxidize any 
iron present to the ferric condition by heating with a little strong nitric acid, 
then add sodium hydroxide solution until a strongly alkaline solution is obtained, 
boil and filter. Examine the filtrate by § 10. ‘« 

The precipitate may contain iron, chromium, uranium and manganese 
(with a little nickel) and the filtrate may contain aluminium and zine. 

5. Dissolve the precipitate in as little hydrochloric acid as possible, dilute 
the solution with hot water and boil several minutes. Add 5 cc. of ammonium 


ANALYSIS OF GROUP III 191 


chloride solution, make barely alkaline with ammonia and filter promptly. The 
precipitate contains all the iron, chromium and uranium; the filtrate, which 
may contain manganese, and traces of nickel, is analyzed by § 9. 

6. Dissolve the precipitate in as little hydrochloric acid as possible, add a 
large excess of ammonium carbonate solution, heat gently but do not boil long, 
and filter. The precipitate contains the iron and chromium; the filtrate 
contains uranium in solution as ammonium. uranyl carbonate. Analyze 
by § 8. 

7. Test the precipitate obtained in (6) for iron by dissolving a part of it in 
a few drops of hydrochloric acid, diluting with a little water and adding potas- 
sium ferrocyanide solution. The formation of Prussian blue shows the presence 
of iron. Test another portion of the precipitate for chromium by mixing it 
with sodium carbonate and potassium nitrate and fusing to form sodium 
chromate. Dissolve the melt, which is yellow if chromium is present, in water, 
acidify the aqueous extract with acetic acid and add a drop of silver nitrate 
solution; a red precipitate of silver chromate is formed if chromium is present. 

8. To test for uranium, add hydrochloric acid to the solution obtained in (6) 
and treat the slightly acid solution with potassium ferrocyanide; a brown pre- 
cipitate shows the presence of uranium. 

9. Test the filtrate from (5) for manganese. Evaporate the solution to 
dryness, dissolve the residue in a little water and add a few drops of potassium 
cyanide solution. Dilute with water, add ammonium sulfide and boil. A 
flesh-colored precipitate is MnS. To confirm the test, dissolve a portion of the 
precipitate in concentrated nitric acid, add a little lead peroxide and boil. 
Dilute with water and allow the precipitate to settle; manganese is shown by 
the characteristic color of the permanganate ion. Or, a small portion of the 
manganous sulfide can be fused with sodium carbonate and potassium nitrate. 
A green melt shown shows the presence of manganese. 

10. Test the filtrate obtained from (4) for aluminium and zinc. Make it 
acid with hydrochloric acid and then add a slight excess of ammonia. A white 
precipitate is Al(OH);. Filter and test the filtrate for zinc (11). To confirm 
the aluminium test, dissolve the precipitate or a small portion of it in nitric acid 
and add half as many drops of 1 per cent cobaltous nitrate solution as there 
are presumable milligrams of aluminium in the precipitate and evaporate the 
- solution nearly to dryness in a casserole. Soak up the solution in a small piece 
of filter paper, roll up the paper and wind a platinum wire around it. Heat in 
the flame till all the paper is consumed and then ignite strongly. Thénard’s 
blue shows the presence of aluminium (p. 1381). 

11. To test for zinc, acidify the solution obtained in (10) with acetic acid and 
saturate it with hydrogen sulfide. Filter and dissolve the precipitate in a little 
nitric acid. Add 1 drop of é¢obalt nitrate solution and as many more drops as 
there are estimated to be centigrams of zinc present. Evaporate to dryness 
and ignite the contents of the dish until the purple color of the cobalt salt disap- 
pears, Rinnmann’s green (p, 188) shows the presence of zinc, 


192 


REACTIONS OF THE METALS 


TaBLE VI.—ANAtysis or Group III in ABSENCE OF PHOSPHATE 
Method B 





Solution may contain Fe++, Fet++, Al+++, Crt++, Mn++, Zntt, Cott, Nit+, 
and Groups IV and V (p. 111 or p. 118). 
Add NH,OH and (NH,)S._ Filter and examine filtrate for Groups IV and V. 


Dissolve ppt. in HCl and HNO3. 


Evaporate, treat with NaOH and Na,O2 and — 















































filter. (2). 

Precipitate: Fe(OH);, H2MnO;, Co(OH)s,|| Filtrate: AlO.-, CrOs, 
Ni(OH).[Zn(OH),]. Dissolve in HNO; and H.2O2.|| HZnO.-. Acidify with HNO; 
Evaporate and boil with conc. HNO3;+KCI103. (8) || and add NH,OH,. (12) 

Precipi-|Filtrate: Fet++, Cot+, Nitt, [Znt+]. ||/Precipi-|Filtrate: CrO,, Zn. 

tate:| AddNH,OH. (5) tate :| Add HC.H;0, and 

MnOz AIOH);.| BaClh. (14) 

LEpeoesa Precipi- [Filtrate: Co(NHs)st+t, Ape 

+tate: Ni(NHs).tt, [Zn(NHs3)¢6t+ +]. al . 
+ 202) Pe(OH);.| Saturate with HyS and treat ppt. metho dt Tecipi- ete 
and test . A. (13)} tate: Zn. Sat 
Test for| with 2-normal HCl. (7) BaCrO,: |/sanaee 

for Mn Fe as in Di 3 ith H. 

with | Method|pecdue: ; + eae | 

NaBiOs.| ‘4 (G Residue: |Solution: Zntt, in H Cl and con- 

(4) - (6) CoS,NiS.} traces of Cott and and | firm zine 

Test for| Nit+. Add NaOH H,.SO3.| asin 
Ni and} and Na,O2. (9) Evapo- | Method 
Co as in rate. A. (16) 
Method |py eeipi-|Filtrate: Green 
A. (8) | tat e:| NasZnOp. color 
Co(OH):| Acidify shows 
; Ni(OH)s.| with Cr. (15) 
Add to| HC,H;0, 
residue | and sat- 
of COS,) urate 
NiS.(10)| with HS 
Confirm: 
Zn as in 
Method 
A. (11) 
PROCEDURE : 


1. Precipitate with ammonium sulfide as in Method A, filter and examine 


the filtrate for the alkali and alkaline-earth metals. 


2. Digest the sulfide precipitate with hot, 6-normal hydrochloric acid anit 


add enough nitric acid to dissolve the nickel and cobalt sulfides. 


water and filter off the residual sulfur. 
ness to remove the excess of acid, dilute to about 25 ce. and carefully neutralize 
If a very heavy precipitate is produced, it is 
To the cold solution carefully add a 


with pure sodium hydroxide. 
best to dilute with a little more water. 
little sodium peroxide powder. 


D 


ilute with 


Evaporate the solution nearly to dry- 


(On account of the violent reaction with water, 


and the fact that the powder often contains a little free sodium, care should be 


ANALYSIS OF GROUP III 193 


taken not to add the peroxide too fast or to a hot solution. Only a little perox- 
ide should be taken from the container at one time and it should be transferred 
directly to glass and never to paper.) Finally boil the solution to decompose 
the excess of peroxide, dilute with an equal volume of water and filter. The 
precipitate contains ferric hydroxide, hydrated manganese dioxide, cobaltic hydrox- 
ide and nickelous hydroxide. The filtrate contains sodium aluminate, chromate 
and zincate. ‘The separation is faulty in the case of zine which normally stays 
in solution. “As much as 5 mg. of zinc may be carried down with the precipitate 
when much iron, nickel or cobalt is present, and as much as 20 mgs. by con- 
siderable manganese. This is probably due to the amphoteric nature of the 
precipitated hydroxides and the insolubility of the zinc salts of the corresponding 
acids. Examine the filtrate for chromium, aluminium and zinc by § 12. 

3. Dissolve the precipitate in hot, 6-normal nitric acid, adding as much hydro- 
gen peroxide as is necessary to reduce the manganese and cobalt to the bivalent 
condition (cf. p. 33). Evaporate the solution nearly to dryness, add 15 ce. 
of 16-normal nitric acid and about 1 gm. powdered potassium chlorate and heat 
_ to boiling. Add 10 cc. more of concentrated nitric acid, heat to boiling, remove 
the flame and add 0.5 gm. more of potassium chlorate. Repeat the treatment 
with fresh portions of chlorate until about 3 gms. of chlorate have been used. Do 
not add the chlorate to the nitric acid solution while it is boiling, as an explosion 
is likely to result, but boil after the addition of each portion of chlorate. The 
treatment with chlorate is best accomplished in a 250-cc. Erlenmeyer flask. 
If a precipitate of MnO, is formed, filter through a thin layer of good-quality 
washed asbestos which is supported by a little glass wool in an ordinary funnel. 
Test the filtrate for manganese by adding 1 gm. of potassium chlorate and boil- 
ing again. Wash the precipitate with a little concentrated nitric acid, which 
has been freed from nitrous acid by boiling with a little potassium chlorate 
just previous to use. Examine the filtrate for iron, cobalt, nickel and zine 
by § 5. 

4. Dissolve the precipitated manganese dioxide with a little hot 6-normal 
nitric acid and a few drops of hydrogen peroxide. Boil to decompose any excess 
of the latter and cool to room temperature. Add a little solid sodium bismuth- 
ate, shake and let the solid settle. A purple solution shows the presence of 
manganese (cf. p. 167). 

5. Add an excess of ammonia to the filtrate from (3) to precipitate ferric 
hydroxide, leaving cobalt, nickel and possibly some zinc in solution as soluble 
complex metal-ammonia cations. Examine the filtrate by § 7. 

6. Examine the precipitate for zron as in Method A. 

; 7. Saturate the filtrate from (5) with hydrogen sulfide, filter off the precipi- 

tated cobalt, nickel and zine sulfides and reject the filtrate. Digest the pre- 
cipitated sulfides with 2-normal hydrochloric acid to dissolve any zine sulfide 
that may be present; a little nickel and a trace of cobalt may be dissolved by’ 
this treatment. Examine the solution for zine by § 9. 

8. Test the sulfide residue for nickel and cobalt as in Method A. 

9. Neutralize the solution obtained in (7) with sodium hydroxide and add 
sodium peroxide as in (2). Examine the filtrate by § 11. 

10. If deemed advisable, any precipitated Ni(OH). and Co(OH); may be 
added to the sulfide residue obtained in (7) and tested for nickel and cobalt as 
in (8). 

11. Acidify the filtrate from (9) with acetic acid and saturate the solution 


194 REACTIONS OF THE METALS 


with hydrogen sulfide. Any precipitate that forms is probably zine sulfide. 
Confirm the zinc test as in Method A. 

12. Acidify the filtrate from (3) with nitric acid and add ammonia until 
present in slight excess. Heat to boiling to coagulate any precipitated alumin- 
ium hydroxide and filter. Test the filtrate by § 14. 

13. Confirm the presence of aluminium as in Method A. 

14. If chromium is present in the filtrate from (12) it is shown by the 
yellow color of the chromate ions. If the solution is colorless at this point it 
is unnecessary to test for chromium. If it is yellow, carefully neutralize with 
acetic acid until a slight excess is present, heat to boiling and precipitate by 
the gradual addition of hot barium chloride solution. Filter and test the 
filtrate for zinc by § 16. ; 

15. The yellow precipitate of barium chromate is conclusive evidence of 
the presence of chromium. Sometimes the yellow precipitate is obscured by a 
white precipitate of barium sulfate. To confirm the chromium test, dissolve the 
precipitate of barium chromate by pouring a mixture of 3 cc. 6-normal hydro- 
chloric acid and 10 ec. of saturated sulfur dioxide solution through the filter 
several times. LEvaporate the filtrate to dryness in a porcelain dish, taking 
care not to overheat the residue, and add a few drops of water. A green color 
shows the presence of chromic ions which were formed by the reduction of the 
chromate. Sometimes a yellow color is obtained during the evaporation. This 
is due to the presence of a very little ferric chloride which has gotten into the 
solution accidentally. It does not interfere seriously with the test. 

16. Saturate the filtrate from (14) with hydrogen sulfide and if a pe 
of zine sulfide is formed confirm the test as in Method A. 


METALS OF GROUP II. HYDROGEN SULFIDE 
GROUP 


MERCURY, LEAD, COPPER, BISMUTH, CADMIUM, ARSENIC, ANTIMONY, 
TIN (GOLD, PLATINUM’ 


MERCURY, Hg. At. Wt. 200.6 
Sp. Gr.=13.60. M. Pt.=—38.7°. B. Pt.=357° 


Occurrence-—Mercury occurs in nature chiefly in the form of 
rhombohedral cinnabar, HgS; from the ore, free mercury is obtained 
by sublimation. According to G. F. Becker,* cinnabar is deposited 
from solutions of its thio salt. The richest deposits are those of New 
Almaden in California, where it occurs with serpentine; of Almaden 
in Spain, Idria in Carniola, and Moschellandsberg in the Palatinate of 
the Rhine. With cinnabar small quantities of native mercury are 
often found. Mercury is also an important constituent of many varie- 
ties of tetrahedrite. 

Metallic mercury is the only one of the metals which is liquid at 
ordinary temperatures. It is insoluble in hydrochloric and dilute sul- 
furic acids, but is soluble in hot concentrated sulfuric acid with evolu- 
tion of sulfur dioxide, forming mercurous or mercuric sulfate accord- 
ing to whether the metal or the acid is present in excess: 


Hg+2H2804 = HgSO4+2H20+S80z2 fT 
QH¢+2H2SO04 = Hge8O1+2H20+802 ft 


Hydrobromie acid hardly attacks the metal at all, while in hydriodic 
acid the metal dissolves readily with evolution of hydrogen: 


Hg+4HI = H2[Hgl4]+He Tt 


The position of mercury in the electromotive series (p. 41) shows 
that mercury cannot be oxidized by H* except when the concentration 
of Hg**t is extremely low. This explains why mercury does not dissolve 
in dilute hydrochloric or sulfuric acid. It seems remarkable, therefore, 
that mercury should be oxidized by hydriodic acid. The reason the 





* Geology of the Quicksilver Deposits of the Pacific Slope. Washington, 1888. 
195 


196 HYDROGEN SULFIDE GROUP 


hydrogen of hydriodic acid can accomplish the oxidation of the mercury 
is because the compound H2{HglI4] is scarcely dissociated at all into 
Hg** ions (ef. p. 10). 

The proper solvent for mercury is nitric acid. 

If the metal is treated with hot concentrated nitric acid, mercuric 
nitrate is formed: 


3Hg+8HNO3=3Hg(NO3)2+4H20+2NO fT . 


If, however, cold nitric acid is allowed to act upon an excess of mer-: 
cury in the cold, mercurous nitrate is obtained: 


Hg(NOs) 2 +Hg = Hge (NOs) 2. 


Mercury is attacked by chlorine, forming calomel (mercurous 
chloride) : 
2H¢g+Cle = HgeCle. 


Two oxides of mercury are known: yellow or red mercuric oxide, 
HgO; black mercurous oxide, Hg2O. | 

These oxides are basic anhydrides, from which two series of salts 
are derived: (a) The mercuric salts, which contain Hgt*, and (b) the’ 
mercurous salts, which contain the group Hget*. We will consider 
first the more stable mercuric salts. 


Mercuric Salts 


Mercuric salts are mostly colorless. The iodide is red or yellow. 
By heating the red tetragonal crystals of mercuric iodide a yellow sub- 
limate of orthorhombic needles is obtained, which gradually changes 
back to the red tetragonal modification; very quickly, almost instantly 
if the yellow crystals are rubbed. This is a general property of dimor- 
phous bodies; the more symmetrical form is almost always the more 
stable. 

The sulfide is black or red. 

Mercurie chloride is soluble in water, 100 cc. of water dissolving 
6.57 gm. at 10°, 7.39 gm. at 20 cc., 11.34 gm. at 50 cc., 24.8 gm. at 
80° and 53.96 gm. at 100°. 

In water containing hydrochloric acid, mercuric chloride is much 
more soluble than in pure water; and in fact the solubility increases 
with the concentration of the hydrochloric acid, due to the formation 
of the complex acid H2[HgCl4]. Alkali chlorides also help ‘to dissolve 
mercuric chloride, forming salts of this complex acid. Mercuric chloride 
is more soluble in alcohol and in ether than it is in water. 

The aqueous solution of mercuric chloride is a poor conductor of 


MERCURY 197 


electricity; it is dissociated to a slight extent only and acts quien 
differently in many cases from a solution of the nitrate, which is a good 
conductor of electricity and therefore contains a good many mercuric 
ions. The cyanide differs from the nitrate even more, as we shall see. 

Mercurie bromide is difficultly soluble in water (94 cc. of water 
at 9° dissolve only 1 gm. of the bromide), but is readily soluble in alcohol, 
and still more soluble in ether. The iodide is more difficultly soluble. 

The halogen compounds of mercury readily form complex com- 
pounds with the halogen compounds of the alkalies, which are very 
stable. 

Mercury compounds are furthermore characterized by the readi- 
ness with which they undergo hydrolysis, forming insoluble basic 
salts. Thus the sulfate is decomposed when diluted largely with 
water (particularly on warming) into a yellow insoluble basic salt: 


3HeSO1+2H20 — 2H2S014-+Hg30280x. 


The presence of hydrogen ions prevents this hydrolysis. 
The nitrate also is readily hydrolyzed into more or less insoluble 
basic salts, according to the dilution. 


Hg(NO3)2+H20 @ HNO3+He(OH)NOs, 
or, | 
2Hg(NO3)2+2H20 — Hg20 (OH)NO3+3HNOs. 


REACTIONS IN THE WET WAY 


A solution of mercuric chloride and one of mercuric nitrate are 
used for these reactions. 
1. Potassium Hydroxide precipitates yellow mercuric oxide: 


HgCle +20H — 2Cl->-+H20 +HgO. 


The hydroxides of the noble metals are exceedingly unstable; 
they lose water, as a rule, even in aqueous solution, forming the anhy- 
drous oxide. 

On adding a lesser amount of caustic potash to a solution of mercuric 
chloride, a reddish-brown precipitate of basic chloride is obtained: 


2HeClo+20H- > 2CI-+H20+Hg20Ch, 
or 
3HeClo-+40H — 4CI-+2H20+Hg302Ch. 


198 HYDROGEN SULFIDE GROUP 


Pe Mercuric oxide and the basic salts are readily soluble in acids. 
2. Ammonia produces ir a solution of mercuric chloride a white 


precipitate of mercurjgaminochloride:’ 
HgCle+2NH3 — NH,zt+ CE +Hg (N H2)Cl. 


t 


This compound, the so-called “ infusible precipitate,” volatilizes 
before it melts. Itris soluble in acids, and in hot ammonium chloride, 
forming the “ fusible precipitate ”’ 


Hg(NH2)Cl+NH4Cl=Hg(NH3)2Cle. 


If ammonia is allowed to act upon mercuric nitrate a white oxy- 
amino compound is always formed: . 


| Tues 
2Hg(NO3)2+4NH3+ H20 — 83NH4NO3+0 NH2-NOs. 


3. Potassium Iodide produces a red precipitate of mercuric iodide, 
HgCle+2I- — 2Cl"+Hglz, 
soluble in excess iodide ions, forming a colorless complex anion: 


Hgl2+21° > [Hela]. 


This complex anion is scarcely dissociated at all into simple mercurie 
cations, for the solution gives no precipitate with caustic soda or potash. 
The alkaline solution is the so-called “‘ Nessler’s reagent,’’ and serves for 
the detection of very slight traces of ammonia. Erte is formed in this 


reaction the brown-colored compound, o< “8 Nv I, which is 
Hee 


soluble in an excess of the “ Nessler’s reagent,’ with an intense yellow 
color (cf. p. 90). 

4. Alkali Carbonates precipitate from both the chloride and the 
nitrate a reddish-brown basic carbonate in the cold, 


AHgClo+4Na2CO03 =8NaCl+3C02 T +Hg.03-COsz, 


which on boiling loses carbon dioxide and is changed into yellow mer- 
curic oxide. 

5. Alkali Bicarbonates produce no precipitation in a solution of 
mercuric chloride, but do cause precipitation from mercuric nitrate: 


4H g(NOz3)2+8NaHCO;=8NaN0O3+4H20+7CO2 T +Hgs03COs. 


6. Hydrogen Sulfide produces in solutions of mercuric salts a 
precipitate which is at first white, then yellow, brown, and finally 


. > 


"* 


MERCURY 499 


black. The white precipitate is formed according to the followin 
equation: 
3HgClo+2H2S > 4HCl+Hg;ClSs. 


By the further action of hydrogen sulfide, black mercuric sulfide is 
finally obtained: 
HgsCloS2+ Hes =2HCl 4 3H es. 


Mercuric sulfide is insoluble in dilute boiling acids. Hot concen- 
trated nitric acid transforms it gradually into white Hg3S2(NOs)2, 


9HgS+8HNO3=2NO 7 +3S+4H20 + 3Hg3S2(NO3)o, 


which by long boiling is changed into the soluble nitrate. 
It dissolves readily in aqua regia, forming the chloride with separa- 
tion of sulfur: 


3HgS+6HCl1+2HNO3=3HgCle+388+2NO 7 +4H20. 


Mercuric sulfide is insoluble in caustic soda and potash solutions, 
and in ammonium sulfide, but it dissolves readily in sodium or potas- 
sium sulfide: 

HgS+K2S = Hg(SK)po. 


By dilution with water this compound is completely hydrolyzed 
into mercuric sulfide, potassium hydrosulfide and potassium hydroxide: 


He(SK)2-+H20 @ KOH-+-KSH+HgS. 


Therefore it is always necessary to dissolve the mercuric sulfide 
with considerable potassium sulfide, or with little potassium sulfide 
and considerable caustic potash, in order to prevent this hydrolysis: 

The fact that Hg(SK)2 is so readily hydrolyzed explains the forma- 
tion of cinnabar in nature. In the interior of the earth the thio 
compound is formed, which is brought by springs to the surface and there 
undergoes the above decomposition. 
7. Potassium Cyanide produces in a solution of ‘mercuric chlo- 
ride no precipitation, because the cyanide, as well as the chlo- 
ride, forms readily soluble complex compounds with alkali chlorides. 
The following are known: K[HgCls], Ko[HgCls], K{Hg(CN)eCl], 
Ke[Hg(CN)2Cle] and Ke[Hg(CN)4]. » 

In a concentrated solution of merctiric nitrate, potassium cyanide 
produces a precipitate of mercuric cyanide, soluble in considerable 
water and jn potassium cyanide: 


Hg(NO3)2+2CN — 2NO3 +Hg(CN)>2. 


200 HYDROGEN SULFIDE GROUP 


Mercuric cyanide is the only cyanide of the heavy metals that is 
soluble in water. It dissolves mercuric oxide perceptibly, forming 
the complex compound (HgCN):O- Mercurie cyanide is not pre-— 
cipitated by alkali carbonates or by caustic alkalies, because the 
mercuric oxide is soluble in mercuric cyanide. It is not decomposed 
by dilute sulfuric acid, although it is by the halogen acids—most 
difficultly by hydrochloric acid, and most readily by hydriodic acid; 
hydrogen sulfide decomposes it with precipitation of mercuric sulfide: 


Hg(CN)2+H2S8 = 2HCN-+H¢s. 


8. Neutral Alkali Chromates precipitate yellow mercuric chromate 
from both the chloride and nitrate solutions. On long standing or 
by boiling, the precipitate becomes red, a basic salt being probably 
formed. 

9. Alkali Dichromates throw down a yellowish-brown precipitate 
from the nitrate solution, but not from the chloride. 

10. Ferrous Sulfate reduces mercuric nitrate on boiling to metallic 
mercury: 

Hg(NO3)2+2Fet*+ — 2Fe* +*+2NO3°+Heg. 


Mercuric chloride and cyanide are not reduced by ferrous sulfate. 
11. Stannous Chloride reduces mercuric salts, at first to insoluble 
mercurous chloride (calomel), 


2HgCle +Snt wae Sn* sail *+HgoClo, 


and by further action to metal, 
Hge2Cle +Snt + = Snt aa ++9Hg¢. 


Metallic mercury separates out in the form of a gray powder. 
By decanting the solution, and boiling the residue with dilute hydro- 
chloric acid, the mercury appears in tiny globules. 

12. Copper, Zinc, and Iron precipitate mercury from solutions 
of its salts: 

HgCle+Fe = Fe* ++2Cl +Hg,* 


HgCle+2Cu = CueClo+Hg. 


On placing a drop of mercury solution (whether of a mercurous or a 
mercuric salt) upon a piece of bright copper-foil, a gray spot is formed, 
which, after being dried, becomes bright as silver on rubbing. 





* This reaction is employed for detecting metallic iron in the presence of FeO. 
If an excess of HgCl, is present HgeCl, is formed (ef. p. 148). 


MERCURY 201 


Mercurous Salts 


The mercurous salts all contain the bivalent mercurous group 
Hg2** and are changed more or less readily into mercuric salts, splitting 
off one atom of mercury from the molecule. Mercurous salts contain- 
ing oxygen, like mercuric salts, are readily hydrolyzed in dilute aqueous 
solutions; thus the nitrate is decomposed according to the equation 


He2(NO3)2+HOH = HNO3+Heg2(OH)NOs. 


Mercurous chloride (calomel) is insoluble in water and hydro- 
chloric acid, but soluble in nitric acid and aqua regia. 


REACTIONS IN THE WET WAY 


1. Caustic Potash precipitates black mercurous oxide: 
Hge (NOsz)2 +20H” — 2NO3 +HeO +Hg20. 


2. Ammonia produces a black precipitate of mercuric amino salt 
with metallic mercury: 


Hg. . 
4Hg(NOs)2-+4NHs-+H20=3NHsNO2+0C _DNHLNO: + 2H. 
H 


It can easily be shown that this precipitate contains metallic mer- 
cury by rubbing a piece of pure gold over it; silver-lustrous gold amal- 
gam will be formed. 

Mercurous chloride gives with ammonia a mercuric amine. with 
separation of metallic mercury: 


HgeCle + 2NH3 = NH4Cl +Hg (NHo2) Cl +Hg. 


By boiling the black precipitate with dilute hydrochloric acid 
or with concentrated ammonium chloride solution, the mercuric amine 
goes into solution, leaving behind drops of mercury. . 
3. Alkali Carbonates give, first, a yellow precipitation of the car- 
bonate, which quickly becomes gray, owing to the formation of mer- 
curic oxide, metallic mercury, end carbon dioxide: 


Hge (NOsz)2 +NaeCO3 = 2NaNO3 +Hge2COs, 
and 
HgeCO3 = HgO+Hg-+-CO2 7 . 


4, Ammonium Carbonate yields the same precipitate as ammonia. 
5. Hydrogen Sulfide immediately throws down a black precipitate 
of mercuric sulfide and mercury (difference from mercuric salts) : 


Hgo(NO3)2+H28 =2HNO3+HgS-+ Hg. 


202 HYDROGEN SULFIDE GROUP 


The black precipitate does not dissolve completely in potassium sul- 
fide, the mercury remaining insoluble, but in alkali polysulfides it 
dissolves. 

6. Hydrochloric Acid and Soluble Chlorides precipitate white 
mercurous chloride (calomel), 


Hge (NO3)2+2Cl — 2NO37° + Hg2Clo, 


insoluble in water and dilute acids, soluble in strong nitric acid and 
aqua regia. On boiling for a long time with water, calomel becomes 
gray, owing to a partial decomposition into mercuric chloride and 
mercury. 

On boiling with concentrated sulfuric acid, mercuric sulfate is 
formed with evolution of sulfur dioxide and hydrochloric acid: 


(a) Hge2Cl> +He2SO4 = 2HCl+ HgeSO4. 
(b) HgeSO4+2H2804 = 2H20+S802 fT +2HgSOx. 


7. Neutral Potassium Chromate precipitates red mercurous chromate 
on boiling (cf. p. 141): 
Hge(NOz)2+ KeCrO4 = 2KNO3+HegeCrO4. 
8. Potassium Iodide precipitates green cndctneees iodide, 
Hge(NOz)2+2KI =2KNO3+Hgols, 


partly soluble in an excess of the Iprecipitant, with the formation of 
potassium mercuric iodide and separation of mercury: 


HgeIo+21 — [Hgl4]"+Hg. 


9. Potassium Cyanide precipitates metallic mercury, mercuric 
cyanide being formed at the same time: 


Hge(NOs)2+2KCN =2KNO3+He(CN)o+Heg. 
10. Stannous Chloride precipitates gray metallic mercury: 
Hge(NO3)2+Snt * > Sn*+++*++2Hg¢. 


REACTIONS OF MERCURY IN THE DRY WAY 


Almost all mercury compounds sublime on being heated in the 
closed tube. Mercuric chloride melts first, then vaporizes, forming a 
crystalline deposit on the cold sides of the tube. Mercurous chloride 
sublimes; the subliraate is almost white, but there is a slight grayish 
tint owing to the decomposition of a small part of the substance into 


MERCURY | 203 


mercuric chloride and mercury. Mercurie iodide yields a yellow sub- 
limate, which becomes red on being rubbed with a glass rod. Mercury 
compounds containing oxygen (all more or less unstable) yield mercury. 

The sulfide gives a black sublimate. 

All compounds of mercury, when mixed with sodium carbonate 
and heated in a closed tube, yield a gray mirror, consisting of small 
globules of mercury. In order to make the drops more apparent, 
place a piece of filter-paper over a glass rod, and rub the mirror 
with it. The small drops thenj run together into large ones, stick 
to the paper, and can be removed from the glass. 


Detection of Mercury in Urine* 


Treat 500 to 1000 cc. of urine in a beaker with 0.5 per cent hydrochloric or 
sulfuric acid, add 0.5 gm. of brass wool (such as is used for the ornamentation 
of Christmas trees) and, while heating on the water-bath to 60° or 80°, pass air 
through the liquid for from ten to fifteen minutes, to keep it in constant motion. 
Any mercury present is replaced by copper: HgCl.+Cu =CuCl.+Hg. 

The mercury, as fast as it is set free, amalgamates with the excess of copper 
present. Pour off the liquid from the tiny threads of brass and wash thoroughly 
by decantation-with distilled water, then with alcohol, and finally with ether. 
Press the brass thread between layers of filter-paper, to free them from any 
adhering ether, and roll them between the fingers into a small pellet. Intro- 
duce this pellet into a thoroughly cleansed and perfectly dry glass tube, 10 em. 
long, 0.5 em. wide and closed at one end. With the aid of the blast flame, 
draw out a capillary of about 1 mm. width in the tube about 0.5 cm. away from 
the brass, toward the open end of the tube. After cooling the tube, heat the 
bottom of it, in which the sample rests, to dark redness. This causes the mer- 
cury to distill off and it is condensed in the colder portion of the tube in a 

gray mirror consisting of tiny drops. If any considerable amount of mercury is 
present, e.g., more than 1 mg., the drops of mercury ean be distinctly seen with 
alens. If less than 1 mg. of mercury is present in the urine, it is very difficult 
to distinguish the mirror. In this case to make it perceptible, transform the 
mercury into scarlet-red mercuric iodide. To accomplish this, place a small 
crystal or two of iodine in a test-tube and cut off the tube containing the mer- 
cury mirror just above the place where the ball of brass rests and place the 
part of the tube containing the mercury in the test-tube. Cautiously heat the 
bottom of the test-tube over a gas flame. As soon as the violet vapors of iodine 
reach the place where the mercvry was deposited, the latter is transformed, by 
very gentle heating, into the red iodide, which can be seen most distinctly by 
removing the little tube and laying it upon a piece of white paper. This method 
is very sensitive and permits the positive recognition of as little as 0.4 mg. of 
HgCl..f i 





*P. Fisrinaer, Z. anal. Chem. (1888), 27, 526. 

+ For other methods of detecting mercury in urine, see JOLLES, Z. anal. Chem., 
39, 230 (1900), Merart, J. Pharm. Chim. [5] 19, 444 (1889); and OprpENHEIM, 
Z. anal. Chem., 42 (1903), 431. 


204. HYDROGEN SULFIDE GROUP 


Detection of Mercury Vapors in the Air 


Place a piece of pure gold leaf in a small glass tube and draw the air to be 
tested through the tube for an hour, at a rate not greater than one liter per 
minute. Meanwhile evacuate a Geissler tube, of the form shown in Fig. 19, 
by means of a water pump (not a mercury. pump!) and finally close both stop- 
cocks. Place the gold leaf, which now contains as 
amalgam any mercury that was present in the air 
tested, in the tube » and suddenly open the cock a, 
which has a wide bore; this causes the gold to be 
sucked into the tube, stopping at c, the mouth of the 
capillary opening. The next step is to replace the air 
in the tube by hydrogen. Introduce hydrogen gas, 
obtained from a Kipp generator and dried by concen- 
trated sulfuric acid, at a@ and allow the gas to pass out 
at b. After a rapid stream of the gas has passed 
through the tube for three minutes, close the cocks a 
and b, without disconnecting the Kipp generator, con- 
nect b with the suction pump* and evacuate the 
apparatus for a minute or two; then close b, and open 
a (which causes more hydrogen to enter the apparatus); 
close a again, open b and once more evacuate the 
apparatus. Repeat this alternate introduction otf 
hydrogen gas and evacuation five or six times. In 
this way the air is entirely replaced by hydrogen. . 
Finally evacuate the tube for five or ten minutes and 
close the cock b. Place the capillary in front of the slit. 
of a spectroscope and allow the secondary current of 
an induction apparatus to pass through the tube. In 
the presence of the merest trace of mercury, the 
characteristic green line 546 uy is distinctly visible in 
the cold, and with somewhat larger amounts of mercury 
the indigo-blue line at 456 uy can be seen. If the wad 
of gold leaf is cautiously warmed with the Bunsea 
flame, the mercury spectrum appears still more 

Fia. 19. sharply. 

Remark. This test is so extremely sensitive that a 
blank test performed in places where work with mercury has been performed 
will often show the presence of this element in the atmosphere. 

If the apparatus has been once used for the detection of mercury it must be 
thoroughly cleansed before it is used again for this purpose. To this end, 
remove the gold and allow aqua regia to remain in the tube for several minutes. 
Draw out the acid and rinse the tube three times with distilled water, once with 
absolute alecoholt and finally dry by passing dry hydrogen through the tube 
for five minutes, while warming it at the same time. Ignite the gold gently to 








* Between the water suction-pump and the Geissler tube, a calcium chloride 
drying tube should be introduced. 

+ All these operations must be carried out in a space where there are positively 
no mercury vapors present in the atmosphere. 


LEAD 205 


distill off any mercury it contains. If now on introducing the gold and evacuat- 
ing the apparatus, the mercury spectrum is no longer visible, the tube is ready 
for a new experiment. 

It may be mentioned that the two platinum wires in the Geissler tube must 
not be provided with aluminium points, because aluminium amalgamates with 
mercury, and when the points are once amalgamated it is impossible to free 
the tube sufficiently from mercury to permit its use for subsequent experiments. 


LEAD, Pb. At. Wt. 207.1 
Sp. Gr. =11.36-11.39. M. Pt.=327.4°. B. Pt.=1600° 


Occurrence.—Galena, PbS, isometric; cerussite, PbCOs3, ortho- 
rhombic and isomorphous with aragonite, CaCOs3; anglesite, PbSOu, 
orthorhombic, isomorphous with anhydrite, CaSO., celestite, SrSOq, 
and barite, BaSO4; pyromorphite, Pbs(PO4)3Cl, hexagonal; mimetite, 
~Pbs(AsO4)3Cl; vanadinite, Pbs(VO4)3Cl. The last three minerals are 
isomorphous and belong to the apatite group. Other minerals which 
may be mentioned are wulfenite, PbMoO,, tetragonal, isomorphous 
with stolzite, PbWO., and the monoclinic crocoite, PbCrO.. 

Lead is a bluish-gray metal. It is attacked by all acids. As, 
however, most lead salts are difficultly soluble in water, it usually 
becomes coated with a layer of salt, which protects it from further action 
of the acid. Thus lead is immediately attacked by dilute sulfuric 
acid according to the equation 


Pb-+H2S01=PbSO04+Hp ¢ . 


But, as lead sulfate is insoluble in dilute sulfuric acid, the reaction 
quickly ceases. Upon this principle rests the use of ‘‘ lead chambers ”’ 
in the manufacture of sulfuric acid, and the use of “ lead pans” for 
the concentration of the dilute “‘ chamber acid.”’ It has, however, been 
found from experience that the sulfuric acid should not be concentrated 
too much in lead pans—stopping when a 78-82 per cent acid is obtained. 
The protecting layer of lead sulfate is soluble in hot concentrated sul- 
furic acid, forming soluble lead bisulfate, PbSO4-+ H2SO4 —> PbH2(SO4)e, 
so that the hot concentrated acid can act on the freshly-exposed surface 
of lead: 
Pb+3H2804 — 2H20+PbHe(SO4)2+S802 TF. 


Lead behaves quite similarly on treatment with hydrochloric acid. 
On the surface a protecting coating of lead chloride is obtained, which 
is soluble in hot concentrated hydrochloric acid, forming HPbCls. 
Lead is soluble, therefore, in concentrated hydrochloric acid. 


Pb+3HCl=HPbCl3+He f . 


206 HYDROGEN SULFIDE GROUP 


Hydrofluorie acid attacks lead similarly, forming a_ protecting 
layer of lead fluoride, which is insoluble in hydrofluoric acid. Con- 
sequently lead retorts can be used for the distillation of hydrofluoric | 
acid and in the preparation of hydrofluoric acid by means of fluorite 
and sulfuric acid. 

Nitric acid is the proper solvent for lead. Lead nitrate is insoluble 
in strong nitric acid, so that lead does not dissolve in concentrated nitrie 
acid; the solution must be sufficiently dilute to prevent the separation 
of the lead nitrate formed. 

Lead forms the following oxides: lead suboxide, Pb2O; lead oxide 
(litharge), PbO; lead sesquioxide, Pb2O3; minium (red lead), Pb3O4; 
lead peroxide, PbOs. oe 

Of these oxides, PbO alone is the anhydride of a base;* from it 
the salts of lead are derived, in which the lead is bivalent. This lead 
monoxide (litharge) is a yellow powder, which melts at about 980° C., 
and solidifies on slow cooling, forming tetragonal crystals (needles). 
It is slightly soluble in water with an alkaline reaction, and is readily 
soluble in dilute nitric acid. : 

Lead suboxide, Pb2O, is formed as a black velvety powder on heat- 
ing the oxalate to about 300° C.: 


2PbC204=3CO2 T +CO FT +Pb20. 


On heating the suboxide in the air, it becomes readily oxidized to 
lead monoxide. 

Lead dioxide, PbO2z, must be considered as the anhycride of ortho- 
plumbic acid, H4PbO,, or metaplumbic acid, H2PbOs, 


OH 
“OH tug 
Pb OH oF Pb=0 , 
OH NOH 


Orthoplumbic acid Metaplumbic acid 


just as $102, SnOz, MnOz, are anhydrides of silicic, stannic, carbonic, 
and manganous acids. The acid H2PbOs3 is formed by the oxidation 
of lead hydroxide, Pb(OH)s, in alkaline solution by means of hypo-_ 
chlorites, chlorine, bromine, hydrogen peroxide, or potassium persul- 


fate: 
Pb(OH)2+20H-+Cle — H20+2Cl-+He2PbOs3. 


The brown metaplumbic acid which separates out goes over at 
100° C. into the anhydride; and the latter on ignition loses oxygen, 
changing into yellow lead monoxide. The other two oxides of lead, 





* Although Pb(C,H;02), is known. 


. LEAD 207 


Pb203 and Pb3O4, may be regarded as salts of the plumbic acids; 

Pb203 as lead metaplumbate, PbPbO3, and Pb304 as the lead ortho- 
plumbate, Pb2PbOx. 

, Pb2O03 is obtained as a yellow precipitate on gently oxidizing an 

alkaline solution of lead monoxide by means of hypochlorites, halo- 

gens, hydrogen peroxide, or persulfates, 


2Pb(OH)2+20H +Cle — 2Cl-+3H20+Pb203, 
and the red minium, Pb30.4, by igniting lead oxide or lead carbonate 
for some time in the air at about 430° C.: 
3PbO0+0=Pb30.. 


Both Pb2O03 and Pb304 behave chemically as salts; for, on treating 
with nitric acid, brown plumbic acid and lead nitrate are formed, which 
corresponds to the action of nitric acid on, say, lead carbonate: 


PbPbO3 +2H+t— Pb+++HoPbOs, 
PbsPbO.+4H*t = 2Pb+ ++ HO +HoPbOs. 


These salt-like oxides * are perfectly analogous to those of mangan- 
ese; on treatment with hydrochloric acid they yield chlorine, the 
plumbic acid, at first set free, behaving like a peroxide: 


PbO2+4HCl=2H20+PbCle +Cle J ; 
Pbe03+6HCl=3H20+2PbCle+ Clo T ; 
Pb304+8HCl=4H20+3PbCle+Cle fF . 


REACTIONS IN THE WET WAY 


Most lead salts are difficultly soluble or insoluble in water; but all 
dissolve in dilute nitric acid, excepting, perhaps, fused lead chro- 
mate, which is very difficultly soluble. 

1. Potassium and Sodium Hydroxides precipitate white lead hy- 


droxide, 
| Pb**+20H- — Pb(OH)s, 


soluble in an excess of the precipitant, forming a plumbite: 
Pb(OH)2+OH — H20+HPbOz. 


Pb(OH)s is also slightly soluble in water which is free from carbonic 
acid. The aqueous solution of lead hydroxide is slightly alkaline. 





* Besides the lead salts of plumbic acid, alkali and alkaline earth salts are known. 


208 HYDROGEN SULFIDE GROUP 


2. Ammonia precipitates the white hydroxide, insoluble in excess 
of the reagent. 

3. Alkali Carbonates precipitate white basic lead carbonate. Alkali 
bicarbonates precipitate the normal carbonate. 

4. Sodium Phosphate precipitates white lead phosphate, 


38Pb**+4HPO4" — 2H2PO4-+Pb3(PO4)o, 


insoluble in acetic acid, readily soluble in nitric acid, caustie soda 
or potash. 

5. Potassium Cyanide precipitates white lead cyanide, insoluble 
in an excess. 

6. Hydrochloric Acid or Soluble Chlorides precipitate from moder- 
ately concentrated solutions flocculent, white lead chloride: 


Pb*++2Cl- > PbCh, 


difficultly soluble in cold water (135 parts of water dissolve 1 part 
of PbCl2), but much more soluble in hot water; .on cooling the solution, 
lead chloride separates in the form of glistening needles or plates. 
Lead chloride is much more soluble in concentrated hydrochloric acid 
and in a concentrated solution of a chloride of an alkali than it is in 
water, as it forms complex compounds with these substances such as 
HPbCls, KPbCl3, which are, however, decomposed on dilution with 
water, with separation of lead chloride. 
7. Potassium Iodide precipitates yellow lead iodide: 


Pb*t+2I- —> Pblo. 


The iodide is much less soluble in water than the chloride (195 ce. 
of boiling water dissolve only 1 gm. of lead iodide), forming a colorless 
solution from which lead iodide separates on cooling, in*the form of 
gold-yellow plates. 

The iodide dissolves to a considerable extent in hydriodie acid, 
and in a solution of an alkali iodide, forming lead hydriodie acid; 
HPblIs or one of its salts (such as KPbI3), all of which are decomposed 
on dilution, with deposition of lead iodide. 7 

8. Alkali Chromates produce a yellow precipitate of lead chromate, 


Pbtt+Cr0O)>:* PbCrO,4 
and 


2Pb* +4 Cr.07-+2C2H302 +H20 — 2HC2H302+2PbCrO,. 


Lead chromate is insoluble in acetic acid, but soluble in nitrie acid 
and in caustic alkali. 


9. Hydrogen Sulfide produces in dilute lead solutions (from 


LEAD © 209 


slightly acid solutions, as well as from neutral or alkaline ones) a black 
precipitate of lead sulfide: 


Pb*t++H2S = 2Ht-+PbS. 


From hydrochloric acid solutions an orange-red precipitate of lead 
sulfochloride is at first obtained, 


2PbCle + HS aR 2HCl a Pb 2CloS, 


which is decomposed immediately by more hydrogen sulfide, forming 
the black lead sulfide. In this respect lead salts are similar to mercuric 
salts (see p. 198). 

Lead sulfide is soluble in dilute, boiling, 2-normal nitric acid, form- 
ing lead nitrate, with separation of sulfur: : 


3PbS+2HNO3+6H*— 3Pbtt+4H2,0+2NO0-+38S. 


The reaction usually goes a little further; some of the sulfur is 
oxidized to sulfuric acid, forming insoluble lead sulfate. ‘The amount 
of sulfuric acid formed (and therefore of the lead sulfate also) increases 
with the concentration of the acid. 

Lead sulfide is also soluble in strong hydrochloric acid: 


PbS+2H*t = Pbht*++H.S 7. 


10. Sulfuric Acid and Soluble Sulfates cause in solutions of lead 
salts the separation of white, difficultly soluble lead sulfate: 


Pb* i HeSO4 — 2H +PbSOxz. 


One part of the salt dissolves at the ordinary temperature in 22,800 ce. 
of water; in water containing a little sulfuric acid it is still less soluble, 
while in alcohol it is insoluble. Lead sulfate dissolves perceptibly 
in hot, concentrated acids, forming Pb(HSO4)2. On cooling the hydro- 
chlorie acid solution, lead chloride separates out in needles. Almost 
all the sulfuric acid of commerce contains some dissolved lead sul- 
fate. In order to detect this, 200-300 cc. of the concentrated acid 
should be diluted with an equal volume of water and allowed to stand 
twelve hours, whereby the lead sulfate separates as a white powder. 

Besides bhing soluble in acids, lead sulfate is easily soluble in 
eaustic alkalies, and in solutions containing the ammonium salts 
of many organic acids. This last property is of great importance for 
the analytical chemist, as it offers a means for separating lead sulfate 
from barium sulfate, silica, etc., which remain undissolved. Ammo- 
nium acetate or ammonium tartrate is usually used as the solvent. 

The reason lead sulfate dissolves in a concentrated solution of 


210 HYDROGEN SULFIDE GROUP 


ammonium acetate is due to the formation of lead acetate, which is 
ionized to only a very slight extent in the presence of an excess of acetate 
ions: 

PbSO4+2C2H302, — Pb(C2H302)2+804> +2H20. 


From this solution the lead can be precipitated as chromate by the 
addition of potassium chromate, as sulfate upon the addition of dilute 
sulfuric acid or as sulfide by ammonium sulfide. 3 

Similarly, ammonium tartrate dissolves lead sulfate by forming 
a tartrate which does not ionize to any extent into simple lead cations. 


REACTIONS IN THE DRY WAY 


Heated with sodium carbonate on charcoal, all lead compounds 
yield a malleable button, surrounded with an incrustation of the 
yellow oxide. On the oesctal stick also, the malleable button is 
readily obtained. 

Lead glass turns black on heating in the reducing flame, owing 
to the separation of lead. 


BISMUTH, Bi. At. Wt. 208.0 
Sp. Gr.=9.8. M. Pt.=270°. B. Pt. about 1435°. 


Occurrence.—Bismuth usually occurs native with nickel and co- 
balt ores. The following ores are of no great importance: Bis- 
mite, BigOs; bismuthinite, BisS3; emplectite, BizSsCuz; bismutite, 
3[CO3][BiOH]-5Bi(OH)s. 

Bismuth is a brittle, reddish-white metal which crystallizes in the 
hexagonal system. The proper solvent for bismuth (as is the case 
with most other metals) is nitric acid. Hydrochloric acid does not 
attack bismuth, and sulfuric acid dissolves it only on warming. 

Bismuth forms three oxides: bismuth trioxide, BizO3, bismuth 
tetroxide, BizO4, and bismuth peroxide, BizQ5/ | 

Bismuth trioxide is a basic anhydride,* from which the salts are 
derived. Bismuth pentoxide, a brownish substance, acts as an acid 
anhydride, forming an acid, HBiOs, corresponding to metaphosphoric 
acid. Salts of this acid have never been prepared in a pure state. 
On igniting, BizOs loses oxygen, forming yellow BisQ3. It dissolves 
in hydrochloric acid with evolution of chlorine, forming a salt of 
trivalent bismuth: 


Bi.0;+10HCl=5H20+2BiCls+2C lo. 


* Bismuth trioxide acts as a weak acid under some circumstances (ef. foot-note, 
p. 211}, 





BISMUTH 211 


Bismuth tetroxide is a brown powder which is sometimes used as an 
efficient oxidizing agent. Commercial sodium bismuthate is probably 
a mixture of NaBiO3 and BigQOx. . 

Bismuth salts are mostly colorless, and are all insoluble in consider- 
able water, on account of being hydrolyzed into an insoluble basic 
salt; thus the chloride is CuarrtAuyely decomposed into bismuth 
oxychloride, 

BiCls3 +H20 = 2HCI1+BiOCl, 


insoluble in tartaric acid (difference from antimony). 

Bismuth oxychloride is readily soluble in hydrochloric acid, the 
above equation taking place from right to left. The reaction, there- 
fore, is reversible and the relative amounts of water and hydrochloric 
acid present determine in which direction the reaction will go. On 
adding water to a slightly acid solution of BiCl3, a white precipitate 
of the oxychloride appears immediately. On carefully adding hydro- 
_ chloric acid, the precipitate again dissolves, but may be reprecipitated 
by the addition of more water. All the other compounds of bismuth 
act as the chloride. The nitrate yields, at first, an amorphous pre- 
cipitate of BiONOs, 


Bi(NO3)3+H20 @ 2HNO3+BiO(NOs3), 
which becomes more basic on further addition of water, and crystalline: 
2BiO(NO3) +H20 = BizO2(OH) (NOs) +HNOs. 


This last compound is the bismuth subnitrate which is so much used 
in medicine. 
REACTIONS IN THE WET WAY 


:. Potassium Byarasise precipitates, in the cold, al bismuth 


hydroxide, 
Bit t+++30H- — Bi(OH)s, 


which, on boiling, becomes pale yellow: 
Bi(OH)3 — H20+Bi0O(OH). 


Both of these hydroxides are insoluble in an excess of the pre- 
cipitant,* but are readily soluble in acids. 
On adding to the alkaline solution, in which the hydroxide is sus- 





*In very concentrated KOH, Bi(OH); dissolves on warming. On cooling, 
part of the Bi(OH)s3is precipitated, and on dilution all of it. In this case the 
hydroxide acts as a weak acid, like antimony trioxide, 


212 HYDROGEN SULFIDE GROUP 


pended, chlorine, bromine, hypochlorites, or hydrogen peroxide, the 
white or yellowish precipitate becomes brown, owing to the formation 
of bismuthic acid: ; 


BiO(OH)+20H +Clz2 — H20+2Cl"+HBiO3. 


2. Ammonia precipitates a white basic salt (not the hydroxide), 
the composition of which varies with the concentration and with 
the temperature. 

3. Alkali Carbonates precipitate, according to the temperature 
and concentration, a number of basic carbonates; one of which is 
formed according to the following equation: 


2Bit+++3C037+H20 = 2Bi(OH)(CO3)+COr fF. 


4. Sodium Phosphate precipitates the white, granular phosphate, 
insoluble in dilute nitric acid, difficultly soluble in hydrochloric: 


QHPO,4-+ Bit ty HePO4 +BiPOx,. 


5. Potassium Cyanide precipitates the white hydroxide (not the 
cyanide). The cyanide is at first formed, but is hydrolyzed: 


(a) Bit+*++38CN7- — Bi(CN)s. 
(b) Bi(CN)3+3HOH =3HCN-+Bi(OH)s. 


6. Potassium Dichromate added in excess precipitates yellow bis- 
muthyl dichromate, 


Cr2077-++2Bit++4.2H20 2 4H*+ (BiO)2Cr20z, 


soluble in mineral acids, insoluble in caustic alkalies (difference from 
lead). 


7. Hydrogen Sulfide precipitates brown bismuth sulfide, 
2Bit +++3H2S — BioS3+6Ht, 


insoluble in cold dilute mineral acids and alkaline sulfides, soluble 
in hot dilute nitric acid, and in boiling, concentrated hydrochloric 
acid. 

8. Alkali Stannites (an alkaline solution of stannous chloride) 
cause a black precipitation of metallic bismuth.* This very sensitive 
reaction is performed as follows: To a few drops of stannous chloride, 
add caustic alkali until the white precipitate at first produced dissolves 
clear. Add this sodium stannite solution to the cold bismuth solution; 





* VaNnINO and TREUBERT, Ber., 1898, 1113. 


BISMUTH 213 


on shaking, a black precipitate immediately appears. The following 
reactions take place in this test: 


Sn? *+20H™ — Sn(OH)>; 
Sn(OH)2+20H — 2H20+Sn02"; 
3Sn02-+2Bitt*+60H- — 3H20+38n037+ 2Bi. 


In making this test, a too concentrated caustic alkali solution should 
be avoided and the solution must be kept cold, otherwise the stannite 
itself may give a black precipitate. 

If too much caustic potash is used, metallic tin will separate out 
(ef. p. 170): 

2Sn02~+H20 — SnO03°-+20H+Sn. 


If too little caustic potash is used, black stannous oxide will be 
thrown down in the cold, after long standing; quickly on boiling: 


| SnO2"+H20 — 20H-+Sn0. 
9. Potassium Iodide precipitates black bismuth iodide, 
Bit*+++3I- — Bils, 
soluble in excess of the reagent, forming a yellow or orange solution: 
Bilgs +I" > [Bila]. 


By diluting this last solution with not too much water, the black 
iodide is reprecipitated, which, on the addition of more water, is changed 
into orange-colored basic iodide: 


Bil3-++H20 = 2HI-+BiOlI. 
10. Metallic Zinc precipitates metallic bismuth: 


2Bit +*4+3Zn > 3Zn +++ QBi. 


12) 
REACTIONS IN THE DRY WAY 


Bismuth salts color the non-luminous flame a pale greenish white. 
Heated with soda on charcoal before the blowpipe, a brittle button 
of the metal is obtained, surrounded by a yellow incrustation of bismuth 
oxide. 

On heating a compound of bismuth in the upper reducing flame 
(p. 68) of the Bunsen burner, the bismuth is reduced to metal, 
which is volatilized and~burnt to oxide in the upper oxidizing flame. 
On: holding a porcelain evaporating-dish (glazed on the outside and 


214 HYDROGEN SULFIDE GROUP 


filled with water) just above the oxidizing flame, a barely visible 
deposit is obtained, which, on being treated with hydriodie acid, 
is changed to scarlet bismuth hydriodic acid: 


Bi2O3+8HI =3Hs0+2H[Bily). 


The hydriodic acid is easiest obtained by moistening a piece of asbestos, 
held in the loop of a platinum wire, in a solution of alcoholic iodine solution 
and then setting fire to the moist asbestos. By holding the burning asbestos 
under the dish, enough hydriodic acid is formed to change the bismuth oxide 
into the red compound. | 


By breathing on this deposit, the color disappears, but reappears 
as soon as the moisture has evaporated. On exposure to fumes of 
ammonia (by blowing the vapors away from the stopper of an ammonia 
bottle) the deposit is colored a beautiful orange, owing to the formation 
of the ammonium salt of the bismuth hydriodic acid, 


H[Bil4] +NHs3 ay NHag|Bil4], 


which also becomes invisible on being breathed upon. 
By moistening this coating with an alkaline solution of stannous 
chloride, black metallic bismuth is deposited. 


COPPER, Cu. At. Wt. 63.57 
Sp. Gr. =9.84, M. Pt.=1080° C. 


Occurrence. Copper occurs as native copper, Cu; cuprite, CusO; 
chalcocite, Cu2S; chalcopyrite, CuFeS2; malachite, Cue(OH)2COz3; 
azurite, Cus(OH)2(COs)2 and atacamite, Cu2z0(OH)Cl- B20. 

Copper is a light red, ductile metal. 

The proper solvent for copper is nitric acid: 


8Cu+8HNO3 — 3Cut *+6NO3 +4H20+2NO f. 


Bright copper is not dissolved by hydrochloric acid alone, but 
in the presence of a weak oxidizing agent, e.g., ferric chloride, the 
solution of the metal is easily effected. Hot hydrobromic acid dissolves — 
it with evolution of hydrogen, forming cuprous hydrobromic acid: 


2Cu+6HBr — Ha[CueBre]+He 7. 


At the beginning of this reaction the solution usually turns dark 
violet on account of the formation of the cupric salt of cuprous hydro-. 
bromic acid, owing to the copper being somewhat oxidized on the sur- 
face. In this case, however, the solution soon becomes colorless, 


COPPER 215° 


owing to the reduction of the cupric salt by metallic copper. On adding 
water to the clear solution cuprous bromide is precipitated: 


{CusBre]~ ~ —> CueBre+4Br-. 


Copper is not attacked by dilute sulfuric acid, but it dissolves in 
hot concentrated sulfuric acid, forming cupric sulfate with evolution 
of sulfur dioxide: 


Cu+2H2504 — CuSO4+H20+S802 fT. 


The behavior of copper toward acids can be understood by reference to 
the electromotive series (p. 41). As copper is below hydrogen in the series it 
can be oxidized by hydrogen ions only when the concentration of cupric ions 
is kept very low (cf. p. 43). Hydrobromic acid dissolves copper because the 
slightly ionized complex is formed. Sulfuric acid dissolves copper by virtue 
of the oxidizing power of the hexavalent sulfur. 


Copper forms two oxides: red cuprous oxide, CueO, and black 
cupric oxide, CuO. 

Both oxides are basic anhydrides, forming cuprous and cupric 
salts. Salts of the cuprous series contain the bivalent cuprous group, 
Cugt+, while proee of tke cupric series contain the simple, bivalent 
copper atom Cut * 


Copper is also known in the trivalent condition.* If a nitric acid solu- 
tion of tellurous acid is evaporated to dryness with a little copper nitrate 
and the residue is treated with KOH solution (1 : 5) it dissolves. If to the clear 
solution 4 to 6 gms. of (NH4)2(SOx)2 are added, little by little, while the solu- 
tion is at the temperature of the water bath, it becomes pink and the tellurium 
is present for the most part as telluric acid but to some extent as the potas- 


oes 
sium salt of telluro-cupric acid : K[O-Cu TeO,]. 
N\oZ 


A. Cuprous Compounds 


The cuprous compounds are extremely unstable, being oxidized 
quickly to cupric compounds. The only known cuprous salts are 
those with the halogens, the very unstable sulfate and the sulfite. 
Cuprous salts are colorless, insoluble in water, but readily soluble in 
concentrated halogen acids, forming colorless solutions. Such solu- 
tions contain the unstable cuprous halogen acids, probably of the 
formula H4[Cuz2X¢6), in which ‘‘ X’’ is either chlorine, bromine, or 
iodine. Salts are known which are derived from these acids, e.g., 
K4[CueCle]. 





* Cf. Moser, Z. anorg. Chem., 54 (1907), 119 and Braunrr and Kuzma, Ber., 
(1907), 3362. 


216 HYDROGEN SULFIDE GROUP 

The cuprous’ halogen acids are changed dark on contact with air. 
The chloride becomes brownish black; the bromide, dark violet; — 
probably due to the formation of cupric salts of the cuprous halogen 
acids. 

The behavior of the cuprous halogen acids toward carbon mon- 
oxide is very important; the latter is readily absorbed, forming an 
unstable compound: 


CueCle+2CO+2H20 @ CuzCle:-2CO:-2H20. 


By boiling the solution the compound is decomposed into cuprous 
chloride and carbon monoxide; cuprous chloride is used in gas analysis 
for the absorption of this gas. 


REACTIONS IN THE WET WAY 


A solution of cuprous chloride in hydrochloric acid should be used, 
which may be prepared as follows: Dissolve 2 gms. of cupric oxide 
in 25 ec. of 6-normal hydrochloric acid, pour the solution into a flask — 
and add 0.58 gm. of copper filings. Place several copper spirals in the 
flask, one end reaching up to its neck, stopper the flask, invert it and — 
let it stand several days. The originally dark solution will gradually 
become colorless, when it is ready to be used for the following reactions: 

1. Potassium Hydroxide produces in the cold a yellow precipitate 
of cuprous hydroxide, , 


Cug* ++20H- — Cu2(OH)s, 


which loses water at the boiling temperature, changing to red cuprous 
oxide: 
Cu2(OH)2 =H20+Cu20. 
2. Hydrogen Sulfide precipitates black cuprous sulfide, 
Cuett+H2S —? 2H*t+Cu.S, 


soluble in warm. dilute nitric acid, forming blue cupric nitrate, with 
separation of sulfur: 


3Cu2S +16HNO3 — 8H20+3S+6Cutt+12NO3-+4NO f. 
3. Potassium Cyanide precipitates white cuprous cyanide, 
Cuz**+2(CN)~ — Cue(CN)o, 
soluble in excess, forming colorless complex cuprocyanide anions: — 


Cu2(CN)2+6(CN)” — [Cu2(CN)g] **. 


COPPER ? 217 


This solution contains no appreciable quantity of cuprous ions, and 
gives no precipitation with potassium hydroxide or hydrogen sulfide. 
It is estimated that in a normal potassium cyanide solution the ratio of 
the concentration of the complex anion to that of simple cuprous ions 
is about 1076: 1. This fact is utilized in the separation of copper 
from cadmium. ; 

In the absence of an excess of CN- ions, however, an appreciable 
ionization takes place: [Cue(CNsg)|==—> Cug**++8CN-, and this ioniza- 
tion increases as the solution is diluted. From the diluted solutions 
the compounds K2Cue(CN)4, K[Cue(CN)s] and finally Cue(CN)e2 are 
obtained, which are less complex in nature. 

All of these compounds, even in the solid state, are decomposed by 
hydrogen sulfide with precipitation of black cuprous sulfide. Con- 
sequently, in order to prevent the precipitation of copper by hydrogen sul- 
fide, considerable potassium cyanide must be added, more than enough to 
form the salt Ke[Cuz(CN)s]. 


B. Cupric Compounds 


Cupric salts are either blue or green in aqueous solution; in the 
anhydrous state they are white or yellow. 

The chloride, nitrate, sulfate and acetate are soluble in water; 
most of the remaining salts are insoluble in water, but readily soluble 
in acids. 


REACTIONS IN THE WET WAY 
A solution of copper sulfate should be used. 


1. Potassium Hydroxide produces in the cold a blue precipitate 
of cupric hydroxide, 


Cut t+20H- — Cu(OH)s, 
which on boiling becomes changed into brownish-black cupric oxide. 


Cu(OH), is slightly amphoteric in nature and dissolves in very concentrated 
KOH or NaOH, particularly on warming, with a blue color. (Cf. p. 180.) 


In the presence of tartaric acid, citric acid, and many other organic hydroxy- 
compounds, cupric hydroxide is not precipitated by the addition of caustic 
alkali, but the solution is colored an intense blue. If this alkaline solution is 
treated with d-glucose, aldehydes, arsenious acid or various other substances 
having a reducing power, yellow cuprous hydroxide is precipitated from the 
warm solution which is changed to red cuprous oxide on boiling. An 
alkaline solution of cupric salt containing tartaric acid is commonly used under 
the name of Fehling’s solution. It may be prepared by mixing together equal 
volumes of a solution containing 34.64 gms. of crystallized copper sulfate in 


218 HYDROGEN SULFIDE GROUP 


500c¢ c. of water with a solution consisting of 173 gms. Rochelle salt and 52 
gms. NaOH in 500 ce. of water. It is best to keep the solutions separate until 
they are to be used. Fehling’s solution is a reagent for many kinds of sugar, 
aldehydes, hydroxylamine, etc. 


2. Ammonia.—On adding ammonia cautiously to the solution 
of a cupric salt, a green, powdery precipitate of a basic salt is obtained, 
which is extremely soluble in an excess of the wisi forming an azure- 
blue solution: 


(a) 2CuSO4+2NH.40H = (NH4)2804+Cue(OH) 2804. 
(b) Cuz2(OH)2S804+(NH4)2804+6N H3 =2 ((Cu(N Hs) 4|SO4 -H20). 


On adding alcohol to the concentrated blue solution, the above 
compound is precipitated as a blue-violet crystalline substance, which 
gradually loses ammonia on being heated, leaving behind the cupric 
salt. On conducting ammonia gas over an anhydrous copper salt, 
the ammonia is eagerly absorbed, with the formation of a complex 
cupric ammonia salt: CuCl2+6NH3=[Cu(NHs3)6]Cle. 

These compounds (which contain as a maximum 6NHg3 to one 
atom of copper) are perfectly analogous to the corresponding com- 
pounds of nickel, cobalt, and zinc. By the precipitation of the ammo- 
niacal solution with alcohol, the compound with 4NHg3 to one atom of 
copper is always obtained. , 

[he ionization of the complex cation, 


[Cu(NH3)4]*t— Cur ++4NHs, 


is slight in the presence of excess ammonia, but much more than that of 
the cuprocyanide ion (p. 217). 

(3) Hydrogen Sulfide precipitates from neutral or very slightly 
acid solutions colloidal, black cupric sulfide; which has a tendency to 
form a colloidal solution (p. 58) and run through the filter: 


Cu**+H28 — 2H*+CuS. 


To prevent the formation of a colloidal solution, the solution must contain 
some electrolyte; the hydrochloric acid present when the precipitation is made 
is usually sufficient. Another difficulty frequently encountered is due to the 
readiness with which a part of the cupric sulfide precipitate is oxidized to sul- 
fate by contact with the air. Thus if a filter containing copper sulfide is 
allowed to stand in the air, a little cupric sulfate is formed which is soluble in 
water. Many cases where the cupric sulfide apparently runs through the 
filter are explained in this way. In filtering a copper sulfide precipitate the rule 
should be never to let the filter drain completely until the filtration and washing 
is over, and the washing should be with dilute Byeey sulfide water, which 
serves to prevent any oxidation. 


COPPER 219 


Copper sulfide is soluble in hot dilute nitric acid, but insoluble 
in boiling dilute sulfuric acid (difference from cadmium); it is soluble 
in potassium cyanide, forming potassium cuprous cyanide. From a 
solution of the latter salt the copper cannot be precipitated by hydrogen 


~ gulfide. 


- Copper sulfide is appreciably soluble in ammonium sulfide, but is 
insoluble in potassium or sodium sulfide* (difference from mercury). 
4. Potassium Cyanide produces, at first, yellow cupric cyanide, 
which immediately loses dicyanogen, forming white cuprous cyanide. 
The latter, as we have already seen, forms soluble potassium cuprous 
cyanide with more potassium cyanide: 


2Cutt+4CN- — 2Cu(CN)o; 
2Cu(CN)2— (CN)2+Cu2(CN)2; 
Cuo(CN)2+6(CN)~ — [Cue(CN)s]|>~. 


On adding sufficient potassium cyanide to the blue ammoniacal 
cupric solution, the complex compound will be decolorized, forming 
potassium cuprocyanide, and the reduction of the cupric salt to cuprous 
condition in ammoniacal solution is accomplished at the expense of 
cyanide ions which are oxidized to cyanate: 


2[Cu(NHs)4]*+-+9(CN)- ) 
| +20H- = [Cus(CNsg)]==+(CNO)~+8NH3+H320. 


Hydrogen sulfide will not precipitate cupric sulfide from the color- 
less solution of potassium ‘cuprocyanide provided sufficient potassium 
cyanide is present (difference from cadmium). Sometimes, when 
considerable copper salt is present, the introduction of HeS causes 
the formation of a red crystalline precipitate of hydrorubianic acid, 
(CSNHe)e. Cf. p. 316. 

5. Potassium Thiocyanate, KCNS, precipitates black cupric 
sulfocyanate, 

Cn+++2(CN8)- — Cu(CNS)>o, 


which is gradually changed into white cuprous thiocyanate, or imme- 
diately on adding sulfurous acid: | 


2Cu(CNS)2+S03°+H20 — 2CNS~ +8047 +2H*+Cu2(CNS)>2. 





*In solutions of alkali polysulfides, particularly out of contact with the air, 
cupric sulfide dissolves with the formation of compounds of the type NHa4[CuSq] 
and K[CuS4]. Cf. Hormann and Hécutien, Ber., 36, 3900 (1903), and Bitz 
and Heras, thid., 40, 974 (1907). 


220 HYDROGEN SULFIDE GROUP 


Cuprous thiocyanate is insoluble in water, dilute hydrochloric 
acid, and sulfuric acid. 

6. Alkali Xanthates produce in aotaiteans of cupric salts, at first, — 
a brownish-black precipitate of cupric xanthate, which splits off dixan- : 
thogen, forming finally yellow cuprous xanthate: 


S S 
sche 
NaS-COC2H;+Cut* — 2Nat+Cu(S-COC2Hs)2. 


Sodium xanthate 


I i ] 
2Cu(S:C-OCeHs)2 — Cus(SCOC2H5)2+ (SCOC2H5)2. 


Cuprous xanthate Dixanthogen 
The reagent, sodium xanthate, is readily obtained by mixing car- 
bon disulfide with sodium alcoholate: 


| 
CS2+Na0C2Hs > NaS-C-OC2Hs. 


The alkali xanthates are not used as reagents in testing for copper, but 
cupric salts are used in testing for xanthates. The reaction is made use of in 
the detection of carbon disulfide in gas mixtures; the gases are allowed to 
act upon sodium alcoholate, whereby sodium xanthate is formed if carbon 
disulfide is present, and the solution after neutralizing with acetic acid is 
tested for xanthate by means of a solution of cupric salt. 


7. Potassium Ferrocyanide, K,4{Fe(CN)¢], produces in neutral and > 
acid solutions an amorphous precipitate of reddish-brown cupric ferro- 
cyanide, , 
2Cutt+[Fe(CN)6]~~ — Cuz[Fe(CN)¢], 


insoluble in dilute acids, but soluble in ammonia with a blue color — 
(difference from molybdenum ferrocyanide which dissolves in ammonia, 
forming a yellow solution). It is also decomposed by potassium 
hydroxide: in the cold, light-blue cupric oxide and potassium ferro- 
cyanide are formed, while, on warming, black cupric oxide is obtained 
(difference from uranium, which yields the yellow uranate both with 
ammonia and sodium or potassium hydroxide). 


REACTIONS IN THE DRY WAY 


The borax, or salt of phosphorus, bead is green in the oxidizing 
flame when strongly saturated with the copper salt; blue if containing 
only a small amount. The reducing flame decolorizes the bead unless 
too much copper is present; in such a case it is reddish brown and 
opaque, owing to the separation of copper. ‘Traces of copper may be 


CADMIUM 221 


determined with certainty as follows: To the slightly bluish bead pro- 
duced by the oxidizing flame, add a trace of tin or of a tin compound. 
Heat the bead in the oxidizing flame until the tin has completely dis- 
solved, then bring it slowly into the reducing flame and finally quickly 
remove it. The bead now appears colorless when hot, but ruby-red 
and transparent when cold. If, however, the bead is kept too long in 
the reducing flame, it remains colorless; but the ruby-red color may be 
produced by cautious oxidation. ‘This reaction is very sensitive, and 
can also be used for the detection of tin. 

Heated with charcoal before the blowpipe (or better still with the 
charcoal stick), spongy metal is obtained. 

Copper salts color the flame blue or green. 


CADMIUM, Cd. At. Wt. 112.4 
Sp. Gr.=8.6. M. Pt.=321°. B. Pt.=770°C. 


Occurrence.—Cadmium is usually associated with zinc in its ores. 
_ It is also found as greenockite, CdS, hexagonal; and as the oxide, CdO,* 
isometric. : 

The most important commercial salt is the sulfate, 3CdSO4:-8H20. 
It is not easily recrystallized. To purify the salt, the concentrated 
aqueous solution is treated with alcohol, and the crystals that are 
deposited thereby are filtered, washed with alcohol and dried upon 
blotting paper. 

Cadmium is a silver-white, ductile metal. Heated in the air, it 
burns to brown cadmium oxide. The proper solvent for cadmium | 
is nitric acid. Dilute hydrochloric and sulfuric acids dissolve it but 
slowly, with evolution of hydrogen. Cadmium ‘forms two oxides: 
black cadmium suboxide, Cd2O, and brownish-black cadmium oxide, 
CdO. 

Cadmium suboxide (whose existence is doubted) is formed with 
cadmium oxide in small amounts when the metal is burned in the 
air. It is also said to be formed, like lead suboxide, by gently 
heating the oxalate away from air. There are no cadmium salts 
derived from this oxide. Cadmium forms only one series of salts, in 
which cadmium is bivalent. 

Cadmium salts are mostly colorless, though the sulfide is yellow 
or orange. Most of the salts are insoluble in water, but readily soluble 
in mineral acids. The chloride, nitrate, and sulfate are soluble in 
water. 





* With smithsonite in the zine deposits of Monte Poni, Sardinia. Chem. Zig., 
1901, 561. 


222 HYDROGEN SULFIDE GROUP 


REACTIONS IN THE WET WAY 


1. Potassium Hydroxide precipitates white, amorphous cadmium 
hydroxide, insoluble in an excess of the reagent (difference from zine 
and lead): 

| Cd**+20H7- — Cd(OH)>. 


On gently igniting the hydroxide the brown oxide is obtained, 
which becomes darker on stronger ignition. The ignition of cadmium 
nitrate yields the black crystalline oxide. 

2. Ammonia also precipitates the white hydroxide, soluble in 
excess (difference from lead), forming complex cadmium ammonia 
compounds, as with zinc, nickel, etc. 


Cd(OH)2+4NHs3 — [Cd(NHs3)4]**F. 


In the presence of normal ammonium hydroxide the ratio of the . 
concentration of the complex anion to that of the simple cadmium 
cation is about 107: 1. In pure water, the ionization takes place to 
a much greater extent; by diluting with water and boiling, cadmium 
hydroxide is reprecipitated from the solution of the cadmium ammo- 
nium compound. 

3. Alkali and Ammonium Carbonates precipitate the white basic 
carbonate insoluble in excess. 

4. Potassium Cyanide precipitates white, amorplfous cadmium 
cyanide, readily soluble in excess: 


Cd**+2(CN)~ — Cd(CN)po, 
Cd(CN)2+2(CN) — [Cd(CN)a]-. 


From the solution of cadmium potassium cyanide the above-men- 
tioned reagents produce no precipitation. In a normal solution of 
potassium cyanide the concentration of the complex anion to that of 
the simple cadmium cation is 10!%:1. This is evidently a much 
weaker complex than the cuprocyanide anion and for this reason cad- 
mium sulfide, though its solubility product is much larger than that of 
cupric sulfide, is precipitated by hydrogen sulfide (difference from 
copper) : 

[Cd(CN)4|>-+H2S — 2CN-+2HCN+CdS. 


5. Hydrogen Sulfide produces precipitates varying in color 
from a canary-yellow, orange to almost brown, according to the con- 
ditions. In neutral solution, whether hot or cold, light-yellow cad- 
mium sulfide is obtained in a condition hard to filter. From acid 
solutions (containing in 100 cc. from 2 to 10 cc. of conc. H2SO4, or 


ARSENIC 223 


from 2 to 5 cc. of conc. HCl) yellow precipitates which turn orange in 

color are at once thrown down and are easy to filter. The latter pre- 
' cipitates are not pure CdS, but always contain more or less Cd2CleS or 
Cde(SO4)S. For this reason cadmium should not be determined quan- 
titatively as the sulfide. 

Cadmium sulfide is insoluble in alkaline sulfides (difference from 
arsenic), but is soluble in considerable hydrochloric acid, warm dilute 
nitric acid, and hot dilute sulfuric acid (difference from copper). 

6. Ammonium Sulfide produces in ammoniacal solutions colloidal 
cadmium sulfide, which has a tendency to form colloidal solutions and 
pass through the filter. The presence of a concentrated salt solution 
prevents its doing this (cf. p. 218).. 


REACTIONS IN THE DRY WAY 


Cadmium compounds, heated on charcoal with soda, give a brown 
incrustation of cadmium oxide. 

If a compound of cadmium oxide is reduced in the upper reducing 
flame of the Bunsen burner, the cadmium oxide is changed to metal, 
which volatilizes, and, in the upper oxidizing flame, goes back to oxide, 
which will be deposited as a brown coating if a glazed porcelain dish 
filled with water is held just above the flame. This oxide always con- 
tains some suboxide mixed with it, and has the property of reducing 
silver oxide to metal; so that if the coating of oxide is moistened with 
silver nitrate solution a black deposit of metallic silver will be obtained: 


Cd.0+2Agt— Cd*+++CdO +2Ag. 


This reaction is very sensitive. 

If it is desired to test the precipitate produced by Ledoiecn sul- 
fide for cadmium in this way, first roast the sample in the oxidizing 
flame and then treat it as just described. 


ARSENIC, As. At. Wt. 74.96 
Sp. Gr. =5.73 


Occurrence.—Arsenic is widely distributed in nature, being found 
in small amounts in almost all sulfides, as, for example, sphalerite and 
pyrites: therefore almost all the zine and sulfuric acid of commerce 
contain arsenic. 

Arsenic occurs native in kidney-shaped masses; also in the form of 
its oxide, As2O3, as isometric arsenolite and orthorhombic claudetite, 
it being dimorphous. 


294 HYDROGEN SULFIDE GROUP 


Mimetite, Pbs(AsO4)3Cl, hexagonal, isomorphous with apatite, 
pyromorphite, and vanadinite, is a well-known mineral containing 
arsenic oxide. 

The most important sources of arsenic are the sulfides, arsenides, 
and sulfo salts: realgar, AseSe, monoclinic; orpiment, As2S3, mono- 
clinic; arsenopyrite, FeAsS, orthorhombic; niccolite, NiAs, hexagonal; 
léllingite, FeAse, orthorhombic; smaltite, (Co,Ni,Fe,)Ase, isometric; 
and proustite, As(SAg)3, rhombohedral. 

Metallic arsenic is a steel-gray, brittle substance. On being heated 
it sublimes, giving off a characteristic garlic odor. The merest trace 
of arsenic may be recognized by this odor. The molecule of arsenic 
contains, like phosphorus, four atoms, (Asg4). : 

Arsenic is insoluble in hydrochloric acid, but readily soluble in 
nitric acid and in aqua regia. 

Dilute nitric acid dissolves arsenic, forming arsenious acid: 


Ass +4HNO3+4H2C = 41H3As03+4NO 7. 


Concentrated nitric acid and aqua regia dissolve it, forming arsenic 
acid: 
3As4 + 20HNO3+8H20 = 12H3As0O4+20NO 7 . 


Arsenic belongs to the same natural group of eldnents as nitrogen 
and phosphorus, and forms, as they do, two oxides, AsgO3 and AseQs5. 


In reality the symbol of the lower oxide is As,Og, but it is customary to use 
the simpler symbol, As2,O3. 


A. Arsenious Compounds 


Arsenic trioxide is formed by the combustion of arsenic in the 
air as white, glistening crystals of regular octahedrons. If the vapors 
of the trioxide are allowed to cool slowly, they solidify to an amorphous 
glass (arsenic glass), which gradually becomes crystalline (white and 
opaque, like porcelain). 

Arsenic trioxide is known in three different modifications: isometric 
arsenic trioxide (white arsenic); monoclinic arsenic trioxide, and 
amorphous, glassy arsenic trioxide. 

The monoclinic modification is difficultly soluble in water (80 ce. 
of cold water dissolve 1 gm. As2O03); while the amorphous, glassy modifi- 
cation is much more soluble (25 ec. of cold water dissolve 1 gm. arsenic 
trioxide). By treatment of the ordinary modification (white arsenic) 
with water, it is not readily wet by the latter; it floats like flour, and 
this behavior is very characteristic. 

The trioxide dissolves quite readily in hydrochloric acid. particularly 


ARSENIC 225 
on warming, from which solution it often separates out, on cooling, in a. 
beautiful, crystalline, anhydrous condition. 
Acting as an acid anhydride it dissolves readily in alkalies, forming 
easily soluble arsenites: 


As203+60H” — 3H20+ 2AsO3 ; 
As203+3C0O3— — 3CO2+2As03°-. 


The tri-metal arsenites derived from the ortho acid H;AsO; are usually 
unstable. Silver arsenite, Ag;AsO;, is the only well-known salt of this type. 
The alkali arsenites are derived from metarsenious acid, HAsO., from pyroar- 
senious acid, H,As.O; or from a polyarsenious acid such as HsAsyOs. The only 
sodium and potassium salts known are of the types KAsO, and K.H,As,QOq; 
of ammonium, (NH,),As,0;. In alkaline solution, however, we may assume 
that AsO;~ ions are present. 


Free arsenious acid, H3AsO3, has never been isolated; as a very 
weak acid it breaks down, like carbonic acid, into water and the anhy- 
dride. 

Arsenic combines with chlorine directly, like phosphorus, form- 
ing the chloride, AsCl3, which behaves exactly like the chloride of 
arsenious acid, similar to PCls. It is a colorless liquid, boiling at 
134° C., and is decomposed quantitatively, like all acid chlorides, 
with water: 

AsCl3+3H20 @ 3HCI+H3AsOs3. 


The aqueous solution of arsenic trichloride, and the solution of the 
trioxide in dilute hydrochloric acid, contain the arsenic as arsenious 
acid. As the concentration of the hydrochloric acid increases, the 
amount of arsenic trichloride increases, until in very concentrated 
hydrochloric acid the arsenic is present almost entirely as trichlo- 
ride. By boiling a solution of arsenic trichloride in hydrochloric 
acid, arsenic trichloride is given off as a gas. If hydrochloric acid 
_is conducted into the solution at the same time (so that the concen- 
tration of the hydrochloric acid is kept as large as possible), all the 
arsenic can be volatilized from the solution as arsenious chloride. 
On evaporating a hydrochloric acid solution of arsenious acid, arsenious 
chloride constantly escapes, so that all the arsenic may be volatilized. 
If, however, the arsenic is present in the form of arsenic acid, no arsenic 
is lost during the evaporation of the solution. 


REACTIONS OF ARSENIOUS ACID IN THE WET WAY 


The arsenites of the alkalies are soluble in water; the remaining 
arsenites are insoluble in water, but soluble in acids. 


226 HYDROGEN SULFIDE GROUP 


1. Hydrogen Sulfide precipitates from acid solutions yellow, floccu- 
lent arsenic trisulfide: 


2As(OH)3+3He2S = 6H20+As28s3, 
2AsCl3-+3H2S =6HCI+As2Ss. 


Arsenious sulfide is insoluble in acids; even boiling 6-normal hydro- 
chloric acid does not dissolve it, but by long boiling with 12-normal 
hydrochloric acid it is slowly changed to volatile AsCls and HS. 
Concentrated nitric acid oxidizes it to arsenic acid and sulfuric acid: 


3As2S3+28HNO3+4H20 =9H2801+28NO J +6F3AsO0u. 
The sulfide dissolves more readily in ammoniacal hydrogen peroxide: 
As2S3+14H202+120H- — 20H20+38047+2AsO4>. 


It is also dissolved by alkalies, ammonium carbonate, and alkali 
sulfides: | 
AseS3+60H- — 3H20+As03 +AsS3 ; 


As2S3+3C037 > 83CO2+As03 ¢AsS3_; 
AsoSs3 +387 <7 2AsS3. 


Just as the anhydride, AseO3, can be referred to the acid, HgAsOs, 
so the thioanhydride, AsgSs, can be referred to the thioarsenious 
acid, Hs3AsS3, which is not capable of existence in the free state, but is 
known in the form of its salts. If one of the latter salts is acidified, 
then thioarsenious acid is set free; but it immediately loses Ho, 
forming the insoluble thioanhydride: 


2AsS3 + 6H* —> 3HeS +AseoS3 is 


On treating a mixture of thioarsenite and arsenite with acid, arsenic 
trisulfide is also precipitated: 


AsO3~ +AsS3 + 6Ht ——y 3H2O +AsoS3 ; 


In this last case precipitation is quantitative only when the solution is 
dilute; from a concentrated solution H.S escapes, so that more HsS must 
be conducted into the solution in order to precipitate all the arsenic. 


This property of forming thio-salts accounts for the fact that 
hydrogen sulfide produces no precipitation from normal arsenites, — 
and only a partial precipitation of AseS3, from mono- and dimetallic 
salts: | 


‘ 


ARSENIC 227 


AsOz3 + 3He2S —> 3H20+AsS3~ rs 
6HAsO37+15H2S — 18H»0+AsoS3-++-4AsS3™ 
3H2AsO3-+6H2S — 8H20+AsoS3-+-AsSa™. 


Consequently, in order to precipitate arsenic completely as tri- 
sulfide, it is always necessary that the solution should contain enough 
free acid to prevent the formation of soluble sulfo-salts. 

2. Silver Nitrate produces in neutral solutions of arsenites a yellow 
precipitate of silver orthoarsenite (difference from arsenic acid), 


AsO3 +3Ag* — Ag3AsOsz, 
soluble in nitric acid and ammonia: 
Ag3AsO3+3H* — 3Ag*+H3As0O3. 
AgsAsO3-++6NH3 — 3Ag(NH3)2+-+AsOs™. 


The first reaction is caused by the formation of nonionized arsenious acid, 
which, in the presence of an excess of H™ ions furnishes even less AsO;= ions 
than are formed by contact of the very slightly soluble Ag;AsO; in contact 
- with water. The solubility in ammonia is due to the fact that the [Ag(NH;).]* 
ion in the presence of an excess of NH; furnishes fewer simple Agt cations 
than Ag;AsO; in contact with water. 


In aqueous solutions of the mono- and dimetallic salts the precipita- 
tion is incomplete: 


OH 
3A OH +3AgNO3=3KNO3+2H3AsO03+Ag3AsO3. 


HeAsO3 +3Ag* = Ag3AsO3+2H". 


In order to make the precipitation quantitative, an alkali (preferably 
ammonia) must be added. As, however, the solution already reacts alkaline, 
it is difficult to reach the exact neutral point. Usually too much ammonia is 
added. As a rule in qualitative analysis it is unnecessary to accomplish complete 
precipitation in this test, as the color of the silver precipitate suffices to show 
whether an arsenite or an arsenate is present. To make the precipitation 
practically complete, add ammonia drop by drop to a solution of silver nitrate 
‘until the precipitate of silver oxide that first forms redissolves; the solution 
then contains complex [Ag(NH;).|* cations instead of simple Agt ions. Add 
this reagent to the arsenite solution which has been made weakly acid with 
nitric acid: | 

H;AsO;+3[Ag(NH;).]*+3H* == 6NH,*+Ag-AsO;. 


The addition of the nitric acid is necessary, as otherwise the solution will 
become ammoniacal, dissolving a part of the silver arsenite. 
In case the solution to be tested contains also a chloride, it should be 


228 HYDROGEN SULFIDE GROUP 


acidified with nitric acid and the chloride precipitated as silver chloride by an 
excess of silver nitrate, and filtered off. To the filtrate, dilute ammonia 
should be added cautiously. At the neutral zone formed by the ammonia 
above the acid solution, a yellow precipitate of silver arsenite will appear. 
This reaction is very sensitive. 


3. Magnesium Ammonium Chloride produces no _ precipitation 
in dilute arsenite solutions in the presence of ammonia (difference 
from arsenic acid). 

4. Iodine Solution is decolorized by arsenious acid, the latter 
being oxidized to arsenic acid: 


H,AsO;~+ I, 2 QHt++ 21-+H:AsO,-. 


To make the reaction take place quantitatively in the direction left to 
right it is necessary to keep the solution neutral; to make the reaction take 
place quantitatively in the direction right to left it is necessary to add a con- 
siderable excess of hydrogen ions. This is in strict accord with the mass- 
action principle. This behavior has been explained by assuming that. free 
hydriodic acid is a better reducing agent than iodide ions, but it is more prob- 
able that the effect of the acid upon the stability of the arsenic compounds is 
more important. In alkaline solutions, the arsenic % more stable in the higher 
state of oxidation and, for this reason, the’ arsenite solutions have strong 
reducing powers in neutral or alkaline solutions. Arsenious acid is ampho- 
teric and forms, as we have seen, a trichloride and trisulfide. In the presence 
of an excess of hydrogen ions from some other source, arsenious acid will 
not ionize appreciably as an acid and the tendency will be to form Ast*t 
cations. Probably arsenic acid is also amphoteric, though to a much less 
extent. In strongly acid solutions, the ionization of the arsenic acid is re- 
pressed to some extent and there is a tendency to form As***** cations, but 
these are far less stable than Ast** cations and, therefore, in strongly acid 
solutions an arsenate acts as a vigorous oxidizing agent. 

The explanation is in the line with the results obtained in the study of 
oxidation potentials (cf. p. 43). The addition of acid decidedly increases 
the oxidizing power of arsenic acid, but slightly diminishes the reduction power 
of an iodide. 

To keep the solution neutral when it is desired to oxidize an arsenite by 
means of iodide, it is not advisable to use caustic alkali solution, as this itself 
reacts with iodine, forming iodide and hypo-iodite. Sodium bicarbonate is 
usually used: 

HCO,-+Ht =e H,O+CoO, T . 


Equally satisfactory is disodium phosphate, which forms with hydrogen ions 
the very slightly ionized H,PO; ions. A normal alkali carbonate can be used 
if the solution is saturated with carbonic acid; there are then not enough OH~ 
ions formed by the hydrolysis of the normal carbonate to react with the iodine. 


5. Stannous Chloride (Bettendorff’s Test)—On adding to con- 
centrated hydrochloric acid a few drops of an arsenite solution and 
then 4 cc. of a saturated solution of stannous chloride in hydrochlorie 


ARSENIC 229 


acid, the solution quickly becomes brown and then black, owing to the 
deposition of metallic arsenic. The reaction takes place more readily 
on warming, but a dilute aqueous solution will not give the reaction. 
In concentrated hydrochloric acid, however, the arsenic is all present as 
trichloride, and this is reduced by the stannous chloride, while arsenious 
acid is not: 

2As*t*++38n*t — 38ntt++++2As, 


B. Compounds of Arsenic Pentoxide 


Arsenic pentoxide, which may be obtained by heating arsenic 
acid, is a white, fusible substance, and is changed by strong ignition 
into arsenic trioxide: AseO5;=As2Os+O2. Arsenic pentoxide is quite 
soluble in water, forming arsenic acid: 


AseOs +3H2O = 2H3AsQOx4. 


Arsenic acid itself may be obtained in the solid state in the form 
of orthorhombic prisms corresponding to the formula 2H3AsO4- HO. 
At 100° C. water escapes, orthoarsenic acid, H3AsO4, being left behind 
as a crystalline powder. 

By gentle ignition more water is given off, forming pyroarsenic acid, 
H4As207, which on further ignition is changed to metarsenic acid, 
HAsOz. In this respect arsenic acid acts exactly like phosphoric acid. 
Both the pyro- and the meta-acids readily take on water, and are 
changed back to the ortho acid. 

The salts of arsenic acid are called arsenates. 

As with orthophosphoric acid, mono-, di-, and trimetallic salts are 
known: NaHeAsO4, NagHAsO,4 and NagAsOu. 

The arsenates of the alkalies are soluble in water; the others are 
insoluble in water but easily soluble in acids. 


REACTIONS IN THE WET WAY 


1. Hydrogen Sulfide on being passed into a cold solution of an 
arsenate in 0.3-normal acid does not cause any precipitation until after 
a long time, when arsenic trisulfide is formed. If the cold solution con- 
tains a large excess of concentrated hydrochloric acid, the arsenic is 
precipitated as pentasulfide. If hydrogen sulfide is passed into a hot 
solution of an arsenate in concentrated hydrochloric acid, a mixture of 
arsenic trisulfide and pentasulfide is formed. 

This behavior is very interesting, but the relations involved are quite com- 


plicated. The solubility products of both arsenic trisulfide and arsenic 
pentasulfide are extremely small, and it requires but a small quantity of either 


230 HYDROGEN SULFIDE GROUP 


Ast*+ or Ast++*+ jons to reach this value even with the sulfur ions from 
slightly ionized hydrogen sulfide. A cold solution of an arsenate in 0.3- . 
normal hydrochloric contains no appreciable quantity of Ast++** cations. 
Arsenic acid is of approximately the same strength as phosphoric acid and it 
is only in the presence of a very large excess of an acid such as hydrochloric 
acid that the ionization of the first hydrogen acid is repressed to a marked 
degree. In the presence of concentrated hydrochloric acid, however, it is 
reasonable to assume that a small quantity of Ast++** cations are present. 
These react with hydrogen sulfide to form the very insoluble pentasulfide, 


H;AsO,+ 5Ht o> 4H.O +Ast fk ws 2Ast +++ +4 5HS — As.S;+ 10H*. 


The arsenic sulfide is so insoluble that the effect of the acid is, on the whole, 
favorable; it favors the formation of Ast+*+** cations and it prevents the 
formation of colloidal solutions of As.Ss. 

Hydrogen sulfide is absorbed by a cold solution of an arsenate in dilute 
acid to a greater extent than can be accounted for by the solubility of hydro- 
gen sulfide in water. Soluble thioarsenates are formed: 


H;AsO.+ H.S —> H;As0O;8 -- H,0. 


Hydrogen sulfide also exerts a reducing effect upon the arsenate. This 
reduction takes place very slowly in the cold, but more rapidly if the tempera- 
ture of the solution is raised or if the concentration of the H is increased (cf. 
p. 228). 


H;AsO,+5Ht+S= > Ast+++4H.0+S8; 2Ast+++3H.S > As.S;+6Ht. 


As soon as the solution contains an appreciable quantity of either Astt+** or 
~Astt* ions, the precipitation of the corresponding sulfide at once takes place. 
The temperature of the solution and the concentration of the acid are exceed- 
ingly important factors in the precipitation of arsenic by means of hydrogen 
sulfide. The pentasulfide is the more insoluble of the two sulfides. 

To precipitate the arsenic quickly by hydrogen sulfide from a solution 
of an arsenate, without employing considerable hydrochloric acid, it is only 
necessary to reduce the arsenic acid by boiling with sulfurous acid, to boil off 
the excess of the latter, and then to conduct hydrogen sulfide into the solution, 
whereby a precipitate of arsenious sulfide is at once formed. 


Arsenic pentasulfide is insoluble in boiling concentrated hydro- 
chloric acid, but, like the trisulfide, it is readily soluble in alkalies, ammo- 
nium carbonate, and alkali sulfides: 


. AsoS;+60H”- =3H20+AsS4 +AsOszS° ; 
As2S5+38CO3- =3COz2 T +AsS4= +As038° ; 
AsoS;+38S~ =2AsS4 . 


By acidifying these solutions, arsenic pentasulfide is reprecipitated: 
2AsS4~ +6Ht aoe. 3He2S +AsoSs, 
AsS4 +As0O3S +6Ht — 3H20+As2S5. 


ARSENIC 231 


Arsenic pentasulfide is oxidized by fuming nitric acid to sulfuric 
and arsenic acids; also by solution in ammoniacal hydrogen peroxide: 


AsoSs5 +20H202+160H =28H20+5804-+ 2AsO04>. 


2. Silver Nitrate precipitates from neutral solutions chocolate- 
brown silver arsenate (difference from arsenious and phosphoric acids): 


AsO4=+3Agt — AgsAsOu, 


soluble in acids and in ammonia. 

3. Magnesium Chloride precipitates, in the presence of ammonia 
and ammonium chloride, a white, crystalline precipitate of magnesium 
ammonium arseniate: 


AsO4=+ Mgt t4NHat— MgNH.AsOxz. 


This precipitate is insoluble in dilute ammonia and is used for the 
quantitative determination of arsenic. By ignition it is changed into 
Magnesium pyroarseniate: 


2MgNH.AsO4 = H20+2NH3+Mg2As207. 


4. Ammonium Molybdate, added in considerable excess to a 
boiling nitric acid solution, precipitates yellow, crystalline ammonium 
arsenomolybdate: 


AsO4~ +3NH4t+ 12M00.47+24H*t— 12H2O+ (NH4)3AsO4-12M003. 


This precipitate, like that of the corresponding molybdenum com- 
pound with phosphoric acid, is insoluble in dilute nitric acid solution 
containing ammonium nitrate, but is readily soluble in ammonia or 
caustic alkali solutions: 


(NH,)3AsO4-12M003+240H~ — 12H20+3NH4*+AsOs- +12M004-, 


The yellow precipitate is also soluble in a solution containing an 
alkali arsenate; complex anions containing more arsenic, are formed 
and the ammonium salts of these complex ions are soluble in nitric 
acid. Consequently a large excess of ammonium molybdate should 
be used if it is desired to precipitate arsenic acid. 


As we shall see later, phosphoric acid behaves similarly toward mag- 
nesium salts and ammonium molybdate. If, therefore, phosphoric acid and 
arsenic acid are both present, it is necessary to precipitate first the arsenic 
with hydrogen sulfide, filter, and oxidize the precipitated arsenic sulfide to 
arsenic acid with fuming nitric acid. In such a solution a precipitate produced 
by means of ammonium molybdate or magnesium chloride must be caused 
by arsenic acid. In the same way a precipitate produced in the filtrate from the 
hydrogen sulfide precipitate must be caused by phosphoric acid. This is safer 


232 HYDROGEN SULFIDE GROUP 


than to depend upon the fact that the ammonium phosphomolybdate forms 
more readily at lower temperatures (60°) than does the corresponding arsenic 
compound. 

5. Potassium Iodide, in a solution strongly acid with hydrochloric 
acid, reduces a solution of an arsenate with liberation of iodine (cf. p. 
228) ; 

HgAsO4+5H*+2I- > Ast +*++4H20-+ Io. 


The reaction takes place quantitatively if the iodine is removed — 
by adding sodium thiosulphate. 


C. Reactions which May be Obtained with All Arsenic Compounds 


1. The Marsh Test for Arsenic.—All compounds containing arsenic 
may be reduced, in acid solution, by means of nascent hydrogen to 
arsine, AsH3: 


As203+6H2 =3H20+2AsHs3 | ; As:03+8H2 = 5H20+2AsHs3 T ; 
AseS3+6He =3HeS T +2AsH3 t ‘ 


The sulfides are reduced very slowly, but the oxides are reduced 
quickly even at ordinary temperatures. To produce nascent hydro- 
gen, zine and sulfuric acid are used. 

This very poisonous arsine possesses a property which enables | 
us to detect with certainty the merest trace of arsenic—as little as 
0.0007 milligram As. By conducting the gas through a heated glass 
tube filled with hydrogen, it is decomposed into hydrogen and metallic 
arsenic; and the latter is deposited as a brownish-black mirror on 
the sides of the glass tube, just beyond the place where it was heated. 

This test is extremely sensitive, and must be made with caution, 
as almost all reagents, especially commercial zinc and sulfuric acid, 
are likely to contain traces of arsenic. Incase these are used with- 
out previous testing, arsenic is likely to be found even although it 
may not have been present in the substance itself. 


The Marsh test is particularly useful for detecting the presence 
of very small amounts of arsenic which could not be found by any 
of the previously mentioned reactions. In cases of poisoning and 
for detecting the presence of arsenic in wall-papers, this test, or a 
modification of it, is always used; we will, therefore, discuss it in 
detail. 

Formation and Properties of Arsine 


(a) Formation.—Arseniuretted hydrogen, or arsine, is produced, as above 
mentioned, by the reduction of compounds containing arsenic with nascent 


ARSENIC 233 


hydrogen. For developing the latter, pure zinc and pure sulfuric acid should 
be used. If other metals and other acids are used (e.g., tin and hydro- 
chloric acid, iron and sulfuric acid), the arsenic compound will be reduced; 
but if iron is used, a part of the arsenic is changed to solid arseniuretted hydro- 
gen, which remains in the flask and consequently escapes detection. If tin 
_ and hydrochloric acid are used, a high temperature is necessary in order to 
accomplish the reduction,* while with zine and sulfuric acid the reaction 
takes place readily at the ordinary temperature. Chemically-pure zinc 
dissolves with difficulty in chemically-pure sulfuric acid, so that it is well to 
activate the zinc by the addition of a little foreign metal. The addition of 
a drop of chloroplatinic acid causes at first a more rapid evolution of hydrogen, 
but the reaction soon slows down and is not accelerated by the addition of 
more choloroplatinic acid. Moreover, the addition of chloroplatinic acid 
has the disadvantage of causing considerable arsenic to be held back by the 
platinum; less than 0.005 mgm. of As.0;, cannot be detected in this way.t 
Much better results are obtained by using an alloy of zinc and platinum. 
Thus F. Hefti t found that zinc alloyed with 10 per cent platinum caused a 
more uniform evolution of hydrogen and that the formation of arsine was 
accelerated, while less arsenic was retained by the platinum. With this alloy 
quantities of As.O; as small as 0.0005 mgm. can be detected with certainty. 
The best activating agent, however, is copper in the form of a zinc-copper 
alloy prepared as follows: Melt 20 gms. of the purest zinc in a small Hessian 
crucible, stir a very little pure copper into the molten zinc with the aid of a 
stick of zinc. Pour the molten metal into water, keeping as much as possible 
of the oxide back in the crucible. With this alloy and 15 per cent sulfuric 
acid, a steady, continuous current of gas is obtained and it is possible to 
detect with certainty as little as 0.00025 mgm. of As.O;. 

Arsenic, arsenious oxide, arsenic pentoxide, and arsenic trisulfide are readily 
reduced in alkaline solution by sodium amalgam, aluminium, or Devarda’s 
alloy and caustic potash, forming arsine. . The reduction takes place quickly, 
and the arsine may be detected by the Gutzeit reaction (cf. p. 238). The 
presence of organic matter in solutioh hinders the reaction; 3 cc. of urine 
in which 1 mgm. of As,O; was dissolved showed no trace of arsine after treat- 
ing for hours with Devarda’s alloy and caustic potash solution. In such 
cases the organic substance must be decomposed before testing for arsenic. 
(Cf. pp. 129 and 152.) 

Arsine is also obtained by dissolving many arsenides in hydrochloric 
or sulfuric acid: 

Zn;As.+6HCl =3ZnCl.+2AsH; T e 


The arsenides of iron are attacked by acids only with difficulty, except 
when an excess of iron is present, when, with the help of the nascent hydrogen, 
they are decomposed, forming solid and gaseous arseniuretted hydrogen. 





*Thus VANINO, working at ordinary temperatures, could not detect less than 
0.002 gm. of As,O; by means of tin and hydrochloric acid, and where chloroplatinic 
acid was added, less than 7; mgm. of As,O; could not be found. Z. angew. Chem., 
1902, 82. 

{+ Bernstern, Inaug.-Dissert., Rostock, 1870. 

t Inaug.-Dissert., Ziirich, 1907. 


234 HYDROGEN SULFIDE GROUP 


Consequently iron-sulfide containing arsenic, on treatment with acids, always 
yields hydrogen sulfide contaminated with arsine.* 

Arsenites can also be reduced to arsine by the action of the electric cur- 
rent.’ It is possible to distinguish between an arsenite and an arsenate in this 
way. 

Certain microbes, namely, Penicilliwm brevicaule, when provided with — 
nutriment containing only traces of arsenic, have the power of forming vola- 
tile arsenic compound of a garlic odor, and ‘this may be used as an extremely ~ 
sensitive test for arsenic. By 

(b) Properties—Arsine is a_ colorless, unpleasant-smelling, extremely 
poisonous gas, which, on being heated away from the air, is decomporrae into 
arsenic and hydrogen: 

4AsH; =Asy+ 6H, T 


By heating in the air, it is oxidized to water and arsenic trioxide. Solid 
iodine changes it to arsenious iodide and hydriodic acid: 


AsH;+3I, = AsI,;+3HI. 


This reaction takes place on conducting arsine over solid iodine. This 
property serves to free hydrogen sulfide from arsine, as hydrogen sulfide does not 
act upon solid iodine, but only upon aqueous iodine solutions. Arsine is 
not attacked by hydrogen sulfide at ordinary temperatures, but at 230° C. 
sulfide of arsenic and hydrogen are formed. 

Arsine is a strong reducing agent: silver salts are reduced to metal (see 
p. 239). 


Directions for Performing the Berzelius-Marsh Test 


The apparatus devised by G. Lockemann,f shown in Fig. 20, may be used 
to advantage. . 

In the flask K, of 100 to 150 cc. capacity, place 3 or 4 gms. of zine alloyed 
with copper (cf. p. 233) and about 20 ¢c. of 4-normal sulfuric acid free from 
arsenic. A steady stream of hydrogen is at once evolved, and in twenty 
minutes the air will be entirely driven out of the apparatus. When, at the end 
of about twenty minutes, the gas escaping at b is found to be pure (by collect- 
ing a little in a small tube and holding it near a flame; it should light without 
a sharp explosion), light the hydrogen at b.{ The flame, should be about 2 or 3 
mm. high and should remain so during the whole of the experiment; if it 
becomes higher, cool the solution in K by placing the flask in cold water, 
and, conversely, if the flame is too low add a little more sulfuric acid or place 
the flask in warm water. ’ 

First of all, test the zinc and sulfuric acid to see that they are free from 





* (Chem. Zentralbl., 1902, I, p. 1245.) 

t Z. angew. Chem., 1905, pp. 427 and 491. 

tA safe way to light the flame is to take the small tube used for testing the gas, 
fill it with gas, light it, and bring it slowly to the end of the tube b. If the gas is 
pure the hydrogen in the small tube will burn quietly for some little time. If 
impure, there will be none left in the small tube after it is exploded and this will 
not light the escaping gas at 0, 


ARSENIC 235. 


arsenic. Heat the hard-glass tube at B just before the restriction in the 


tube, which is 5 mm. long and 1.5-2 mm. wide. If at the end of twenty minutes 
. there is no arsenic mirror formed in this capillary, the ‘Teagents are free from ~ 


arsenic. 

Transfer the sulfuric acid solution to be tested for arsenic, and which must 
be free from organic substances, 
sulfides, chlorides, nitrates, or other 
oxidizing agent, to the graduated 
funnel 7’, and add it little by little 
to the flask K without in any way 
interrupting the current of hydrogen. 
Just before adding the solution to 
the flask, light the two burners at 
A and thereby heat the glass tube 
to dull redness. The gas as it 
escapes from the flask K_ passes \ 
through the drying-tube C con-' 
taining granular calcium chloride, 
and then passes into the tube A, 
where any arsine is quantitatively 
decomposed into arsenic and hydro- 
gen. The arsenic is deposited on 
the cold walls of the capillary. Cool 
the end of the capillary, in order to 
form a sharply defined mirror, by 
winding around it a piece of wicking, as shown in Fig. 20, and allowing 
water to drop upon it from the dish W during the experiment. 

All the arsenic will be deposited at the end of an hour, and by comparing 
the mirror with a series of standards the amount can be estimated accurately 
(see page 238). 

Remark.—If{ the tube A is 
not heated at all, but the gas 
ignited at b as above described, 
the arsenic may be deposited 
upon a cold porcelain dish by 
holding the dish in the flame. 
The deposit is readily soluble 
- in sodium hypochlorite solution 
(difference from antimony). In 
this form the test was used by 
James Marsh in 1836. 

Confirmatory Test.—In_ the 

Fic. 21. small glass tube open at both 

ends (see Fig. 21) the arsenic 

mirror is found. Hold the tube in an inclined position and heat it by 

means of a small flame whereby the arsenic is changed to arsenic trioxide, 

giving off the characteristic garlic odor, which can be detected at the upper 

end of the tube if only ;4;5 of a milligram of arsenic trioxide is formed. 

After the tube is cooled, the arsenic trioxide is to be found at a in the form 

of small glistening octahedrons, which can be seen with the magnifying glass 
or often with the naked eye. 




















236 HYDROGEN SULFIDE GROUP 


These three facts—formation of the mirror, the garlic odor, and the 
octahedrons—suffice to prove the presence of arsenic; but the more proofs we 
have, the more certain we are of the accuracy of the result. If the octa- 
hedrons have been recognized, seal the capillary end of the tube with a flame, 
and introduce 1 to 2 drops of pure, con- 
centrated hydrochloric acid into the tube 
with the help of a dropper, and move the 
tube so that the arsenic trioxide is moist- 
ened by the acid; then add 6 to 10 drops 
of distilled water and pass hydrogen sul- 
fide into the tube, whereby yellow arseni- 
ous sulfide is formed. 

The hydrogen sulfide required may be 
generated from a solution of sodium sulfide 
by allowing it to flow into dilute sulfuric 
| acid, as illustrated in Fig. 22. The upper 
¥ N4cotton Part of the test-tube contains a wad of 
cotton wool, which prevents any of the 
solution in the tube from being mechani- 
cally carried over into the tube containing 
the arsenic. 

As an example of the practical appli- 
cation of this delicate test, we shall de- 
scribe the method to be employed in the 
detection of arsenic in wall-papers, etc. 
The amount of arsenic contained in wall- 
papers is usually so small that weighing 
the mirror produced would not be accu- 
rate.* It is best, therefore, to prepare a 
number of mirrors from known amounts 
of arsenic, to establish a scale for deter- 
mining how much is contained in the given wall-paper or fabric. 

First of all, the arsenic must be extracted completely from the paper, and 
to this end it is necessary to decompose the organic material, which is accom- 
plished preferably as follows: 


Na,$ 






































Decomposition of Organic Material 


Take exactly one square decimeter (100 sq.cm.) of wall paper, roll it into 
a cylinder and push it down into a tube closed at one end, such as is used for 
the Carius determination of the halogens (see Volume II of this book). Add 
2 cc. of pure, fuming sulfuric acid (25 per cent oleum, Kahlbaum) through a 
long-stemmed funnel. Then pour 3 or 4 cc. of fuming nitric acid into a small 
test-tube and carefully allow the latter to slip down the sides of the Carius 
tube so that the two acids-do not come in contact with one another. At 
the open end of the tube draw out the glass to form a strong capillary and 





*In Massachusetts, the law permits the presence of 0.1 mgm. per square meter 
in wall-paper, but only 0.01 mgm. per square meter in wearing apparel. In most 
cases it is merely necessary to determine whether the legal limit is exceeded. 

7 C. R. Sanarer, Amer. Acad. of Arts and Sciences, 26, 24. 


ARSENIC 237 


seal the end as described in Volume II. Cover the tube with asbestos paper 
and slowly heat it inside a strong iron tube, in the furnace used for Carius 
tubes, to a temperature of about 230° and keep it at this temperature for 
an hour and a half. Allow the tube to cool and then withdraw it from the 
iron protective tube, by means of a wire previously fastened to it, until the 
capillary projects a little; heat this with a Bunsen flame. As soon as the glass 
becomes soft the pressure on the inside of the tube blows out a hole through 
which gas escapes until the pressure is the same inside the tube as without. 
Break off the point of the tube and rinse the contents of the tube and of the 
tip into a porcelain evaporating dish. The colorless solution * thus obtained 
will contain all the arsenic in the form of arsenic acid. Evaporate the solu- 
tion till fumes of sulfuric acid are evolved thickly, and then, after cooling, 
add 15 cc. of water and pour the liquid into the funnel 7 of Fig. 20, rinsing 
the dish twice with 3 cc. portions of water. After mixing the liquid in 7 by 
means of a small stirring rod, make a note of the total volume of the 
liquid.| During this operation, the Marsh apparatus should be made ready 
for the test. 

When the apparatus is ready, add a few drops of the well-mixed solution 
through 7 to the reduction flask K. If no mirror appears within three or 
four minutes, add one-eighth to one-quarter of the filtrate little by little; 
and if no mirror appears after five minutes, the whole filtrate. The whole 
filtrate is not added at once, because if too strong a mirror is obtained, it is 
much more difficult to estimate the amount of arsenic present. After twenty- 
five minutes all the arsenic will be deposited if not more than 0.05 mg. of 
arsenic is in solution. If a mirror of sufficient density was cbtained in fifteen 
minutes from only a fraction of the whole solution, no more should be added, 
but the operation should be continued for ten minutes more, the flame 
extinguished, and the tube allowed to cool while hydrogen continues to pass 
through it. The mirror is then compared with the scale, and the remaining 
part of the filtrate is weighed in order to determine how much was used for 
the test. 

If sufficient material is at hand, a duplicate experiment should be made 
with a new tube and a new sample. The results of a few such determinations 
are given in the following table: 





Total Weight 

















2 i i : 

Sees | eet || eee | ee) eee 
100 31.63 31.63 0 0 0 
100 30.11 10.23 0.015 0.044 4.4 
arene cla) ree. ne's 9:87 0.013 0.0399 3.99 
100 28.72 8.32 0.045 0.155 15.5 
Choate SEE eck sales 7.53 0.042 0.163 16.3 

50 30.22 2.64 0.015 0.172 34.4 
ree REMC ey 3.22 0.020 0.187 37.4 











*If the paper contained iron or aluminium compounds, the anhydrous sulfates 
are formed. These dissolve by heating with water. The above method is the 
best and cleanest for decomposing organic material. 

+ Instead of measuring the liquid it may be weighed. 


238 HYDROGEN SULFIDE GROUP 


The comparison of the mirrors is best made in transmitted light. The 
normal mirrors are prepared as follows: Dissolve 1 gm. of pure, sublimed 
arsenic trioxide in a little sodium carbonate solution, acidify with dilute sul- 
furic acid, and dilute to a liter. Take 10 ce. of this solution, of which ‘1 ee. 
contains 1 mg. of As,O;, and again dilute to a liter; in this way a solution 
is obtained of which 1 cc. contains exactly 0.01 mg. of As.O;. Measure out 
1 cc., 2 ec., 3 cc., 4 ec., and 5 cc. of the solution and introduce each portion 
separately into the Marsh apparatus, and prepare the corresponding mirrors 
in different tubes. It is best to prepare two tubes from each amount of arsenic, 
as the mirrors are not always the same. These mirrors may be kept in the dark 
for some time; but on exposure to the light they fade perceptibly. Mirrors 
which are sealed up with hydrogen do not keep as well. 


Testing Urine, Blood, Milk, Beer, etc., for Arsenic 


Evaporate 100 cc. of the liquid in question to dryness in a porcelain dish. 
By means of a spatula introduce the residue as completely as possible into a 
tube, such as is used for the Carius determination (Vol. II), and add 4 ee. 
of a 25 per cent oleum. In order to transfer the rest of the residue to the 
tube, pour 2 cc. of fuming nitric acid into the dish and after wetting all the 
sides of the dish, transfer the acid to’ a small test-tube. Repeat this operation — 
twice more and then allow the small test-tube to slip into the Carius tube. 
Draw out a capillary at the open end of the tube and seal it by the flame. 
Heat the tube in the Carius furnace for one hour at 160°. After cooling, — 
open the point of the tube with the usual precautions (p. 237), and release 
the pressure. Seal the tube again and heat for half an hour to an hour at 
230°. The operation is then continued as described above. 

With urine a somewhat different procedure is followed. Evaporate liquid 
not quite to dryness, but to sirupy consistency; spread the sirup upon a porce- 
lain boat and allow it to slip into the Carius tube. The rest of the process is 
carried out as before. 


2. The Gutzeit Test for Arsenic depends upon the behavior of 
arsine toward a concentrated solution of silver nitrate (1 : 1) (accord- 
ing to Eidenbenz, a crystal of solid silver nitrate should be used). 
The silver nitrate is at first colored yellow and then black, the follow- 
ing reactions taking place: : 


1. 6AgNO3+AsH3 = AsAgs -8AgNO3+3HNOs. 
Yellow 


2. AsAg;-3AgNO3+3HOH = H3As03+3HNO3+6Ag. 





The test is carried out as follows: Place a little of the substance in a 
small test-tube, Fig. 23, add a few grains of zinc and a little dilute sulfuric 
acid, and place a wad of cotton near the top of the tube asa filter. Over the 
mouth of the tube place a piece of filter-paper with a crystal of silver nitrate 
on top. 

If arsenic is present, the silver nitrate is at first turned yellow, but it 
becomes black very quickly. 


ARSENIC 239 


This reaction is often used for quickly testing commercial acid for arsenic, 
but it is not as reliable as the Bettendorff test (p. 228), because phosphine * 
and stibine give a similar reaction with silver nitrate, 
while they are not reduced by stannous chloride. AgNO, 

If arsine is allowed to act upon a dilute solution of Fans \ Paper 
silver nitrate, the yellow compound AsAg;-3AgNO; is 
not formed, for it is immediately decomposed hydro- oh 

@| Cotton 


lytically, according to the equation BUEN 
(aaa) wD 


AsH,+6AgNO;+3HOH =6HNO,+H,AsO;+6Ag. teh 


If the precipitated silver is filtered off, and am- 
monia then poured on top of the filtrate, the neutral 
zone will appear yellow owing to ae formation of 
silver arsenite. 

This reaction never takes place ‘iit quantita- 
_ tively; the deposited silver invariably contains a little 
silver arsenide, Ag;As. 

If the silver solution is made ammoniacal, it is 
true that all the arsine .will be absorbed, but the de- 
posited silver still contains a little arsenic and the 
solution a small quantity of ammonium arsenate. If, Fig. 23. 
however, the ammoniacal solution containing the 
silver nitrate and ammonium. arsenite is heated to boiling, then the arsenite 
is oxidized quantitatively to arsenate with deposition of silver. 


2Agt+AsO;-+OH~ — HAsO,-+ 2Ag. 














Under these conditions eight atoms of silver are deposited from each origi- 
nal molecule of arsine: 


AsH;+8Agt+110H~ — AsO,=+7H,0-+ 8Ag. 


If the deposited silver is filtered off and the filtrate carefully neutralized 
with nitric acid, a brown precipitate of silver arsenate is formed. 


Somewhat less sensitive than the original Gutzeit test, although 
very satisfactory, is the modification recommended by Fliickiger ft 
and Lehmann.§ 


Instead of allowing the arsine to act upon silver nitrate, bring it into contact 
with mercuric chloride || paper, which is turned yellow by a little arsine and 
reddish brown by considerable arsine. {] 





* Commercial zine often contains a small quantity of phosphorus. 

{ Cf. REcKLEBEN, LocKEMANN and Eckuarpt, Z. anal. Chem, 1907, 671. 

t Arch. Pharm. (31, 1889, 27. 

_ § Pharm. Zig. Berlin, 1892, 36. 

|| To prepare the mercuric chloride paper moisten some filter paper with an 
alcoholic solution of mercuric chloride, allow the alcohol to evaporate and repeat 
the process four or five times. 

q Aueust Gorruetr, J. Soc. Chem. Ind., 22, 191 (1903). 


240 HYDROGEN SULFIDE GROUP 


The exact composition of these compounds is not known definitely. 
Possibly As(HgCl); is formed first; and then, by further action of AsHs, 
AsH(HgCl)2 and As.Hg; are formed.* 


AsH3+3HgClp— 3HCl+ As(HgCl)3 
2As(HgCl)3+ AsH3— 3AsH(HgCl)> 
As(HgCl)3+ AsH3 — 3HCl+ AsoHg3. 


These arsenic compounds are characterized by their insolubility in 80 per 
cent alcohol. 

Stibine gives no reaction in this test when little of it is present, but the 
presence of somewhat more of it causes the formation of a brown spot which 
is soluble in alcohol. If, however, arsenic and antimony are both present, 
the former is recognized by cutting out the spot from the rest of the filter- 
paper and placing it in 80 per cent alcohol, whereby the brown spot due to 
the antimony is removed in a short time and the yellow arsenic spot appears 
plainly. When considerable antimony is present, the test fails; a gray-black 
spot is produced which does not disappear on treatment with alcohol. 


3. The Reinsch Test is very easy to make, but it is not as sensitive 
as the tests just mentioned. It depends upon the fact that when 
a strip of polished copper foil is added to a solution of arsenious acid, 
the copper is colored gray owing to the deposition of CusAse on the 
copper. | 

From concentrated solutions the arsenic separates out in the 
cold, but from dilute solutions only on warming. If considerable 
arsenic is present, the gray copper arsenide drops off from the copper. 
Antimony is also precipitated on copper from its solutions, so that 
the deposit must be tested for arsenic in the dry way. Arsenic acid 
is also reduced by copper, but. only on warming. 

The Reinsch test is often used in testing wall-papers for arsenic. 
The pieces of paper are treated with a little hydrochloric acid (1 : 2), 
a piece of copper foil added, and warmed. A gray deposit on the copper 
indicates the presence of arsenic. 

To confirm this test, the piece of gray copper foil is placed in a 
tube of difficultly fusible glass and heated in a stream of hydrogen 
gas; an arsenic mirror is produced which can be tested as described 
on p. 235. | 


Detection of Arsenic in the Human Organs 


To detect small quantities of arsenic present in the organs of persons who 
have probably died from poisoning, it is necessary in the first place to destroy 





* Besides the above compounds, AsH2(HgCl) and As,Hg; are said to exist. 
Cf. Parrueit, Arch. Pharm., 237, 121. 


ARSENIC . 241 


all the organic tissue; this may be accomplished by the action of sulfuric and 
nitric acids.* Place 200 g. of the organ, 200 cc. of concentrated nitric acid and 
5 ec. of 2 per cent permanganate solution in a 2-liter flask and heat until 
foaming ceases. Then transfer tlhe solution to a 1-liter flask, rinsing out 
the larger flask with 100 cc. of concentrated nitric acid and 100 cc. of water. 
Boil four hours, or until the solution is reduced to 80 cc. in volume. Add 100 ce. 
of concentrated sulfuric acid and evaporate till fumes of sulfuric acid are - 


evolved. Cool, add 2 or 3 cc. of concentrated nitric acid and again heat until 


white fumes are evolved and repeat this treatment with nitric acid about six 
times. Then, when all the nitrie acid has been expelled, cool, dilute with 
100 cc. of water and test the solution by the Marsh test. 

To determine the quantity of arsenic present, transfer the solution to a dis- 
tilling flask, add 20 ce. of concentrated sulfuric acid, 50 gms. of sodium chloride, 
1 gm. of potassium chloride and 10 gms. of ferrous sulfate crystals. Distill 
into a flask containing 50 gms. of sodium bicarbonate and 100 cc. of water 
until all the solid bicarbonate dissolves. Titrate the sodium arsenite solution 
thus obtained with iodine as described in Vol. II. A blank experiment should 
be made with all the reagents to make sure that they are free from arsenic. 


REACTIONS IN THE DRY WAY 


Metallic arsenic burns, giving off a garlic odor. Mixed with sodium 


carbonate and heated on charcoal, all arsenic compounds give this 


odor. 

Oxygen compounds of arsenic are easily reduced to metal in the 
upper reducing flame. On holding a porcelain dish (glazed on the 
outside and filled with water) directly over the sample, the arsenic 
vapors are condensed on the dish, forming a brownish-black coat- 
ing which is soluble in sodium hypochlorite solution, disappearing 
instantly, the arsenic being oxidized to arsenic acid: 


Aso+5Na0Cl+3H20 =5NaCl+2H3AsOg. 


If the porcelain dish is not helé closely above the reducing flame, 
but above the upper oxidizing flame, the arsenic vapors are burned 
with a bluish flame to white arsenious oxide which deposits on the 
dish. 

If this deposit is moistened with silver nitrate, and ammonia 
vapors blown upon it, a yellow ccloration due to AgsAsOz is formed, 
which disappears if more ammonia is allowed to act upon it (differ- 
ence from antimony): 





* Cf. Gautier, Bull. soc. chim. Paris, 25, 252 (1875); CurrreNDEN and DoNaLp- 
son, Am. Chem. J., 11, 236 (1880-1); JoacnimoGuti, Arch. exp. Path. Pharm., 78, 
1-16 (1914). 


242 HYDROGEN SULFIDE GROUP 
AseO3 + 6AgNOs3 + 3H2O = 2Ag3AsOz3 + 6HNOs. 


The ammonia serves to neutralize the nitric acid formed by the: 
reaction, but the precipitate dissolves in excess of ammonia as well 
as in nitric acid. 


The microchemical method of Hartwich and Toggenburg* is often useful 
when the arsenic is present as trioxide. Prepare a glass cylinder about 12 mm. 
in diameter and 10 mm. in height and make sure that both the upper and 
lower edges are smooth. Place this cylinder upon a small watch-glass and 
pour into it a little of the substance to be tested, well mixed with ignited sand. 
Cover the cylinder with a flat glass slide about 50 mm. square. Heat the 
watch-glass very carefully with the flame from a small burner; the flame 
should not be over 5mm. high and 30 to 40 mm. below the watch-glass. After 
heating ten or fifteen minutes, allow to cool slowly and finally examine the 
bottom surface of the glass slide. If arsenic is present as trioxide, it will have 
sublimed and the vapors will have condensed, upon the slide for the most part, 
in the form of beautiful octahedrons, the shape of which is very distinct when 
viewed under the microscope. These crystals are visible when only 0.01 mg. 
of arsenic is present. The sublimate may be identified further by the silver 
reaction. | 


ANTIMONY, Sb. At. Wt. 120.2 
Sp. Gr. =6.7-6.8. M. Pt.=630°C. B. Pt.=about 1450° C. 


Occurrence.—Antimony seldom occurs free in nature, although 
large amounts of the metal have been found recently in Australia. 
The most important compounds containing antimony are (as with 
arsenic) the sulfur compounds. Stibnite, SbeS3, orthorhombic, is 
found in Japan in beautiful crystals. The occurrence of kerme- 
site, Sb2OSe, is interesting, as this compound is often met with in 
analysis. 

Of the oxygen compounds the dimorphous antimony trioxide is 
known as isometric senarmontite and orthorhombic valentinite. 
Antimony also occurs in many thio salts, of which the tribasic silver 
thioantimonite, or pyrargyrite, AgsSbSs3, may be mentioned. 

Antimony is a silver-white, brittle metal. It burns readily in 
the air to antimony trioxide. The solvent for antimony is aqua regia, 
by which it is converted into chloride. Nitric acid attacks antimony, 
changing it into Sb2O4, which dissolves slightly in concentrated acid, 
but is insoluble in dilute acid. 





* Schweiz. Wochenschrift fiir Chem. u. Pharm. 1909, No. 52, p. 1. 


ANTIMONY 243 


Antimony forms three oxides: antimony trioxide, Sb2O3; anti- 
mony pentoxide, Sb205; and antimony tetroxide, Sb2O04, which may 
be regarded as antimonous antimonate, and is a very indifferent 
substance chemically. Antimony trioxide as a rule shows basic prop- 
erties, while antimony pentoxide has more the character of an acid 
anhydride. 


A. Compounds of Antimony Trioxide 


By burning the metal in the air, the trioxide is obtained, which 
on stronger ignition in the presence of air is changed to the inert 
Sb20.. 

The trioxide is dissolved by concentrated hydrochloric acid, form- 
- ing antimony trichloride, a compound which (like bismuth chloride) 
is readily changed into a basic salt by the action of water, the decom- 
position of which depends upon the concentration of the reacting sub- 
stances. Thus an oxychloride SbOCl is known, which is formed 
according to the following equation: 


SbCl3 + H20 = 2HCl+SbOCI1. 


In the presence of a large amount of water some oxide is also formed 
with the oxychloride: 


2SbOCI+H20 = 2HCI1+Sb20s. 


A mixture of oxychloride and oxide is known as “ algarot ”’ powder, 
Sb203-2SbOCI. 

By boiling with considerable water the oxide alone is obtained. 

Antimony trioxide forms three hydroxides, which behave as very 
weak acids: orthoantimonous acid, H3SbOs; pyroantimonous acid, 
H4Sb205; and the hypothetical metantimonous acid, HSbOs. 

Salts of the metantimonous acid are known, although the free 
acid itself has never been isolated. On boiling the oxide Sb2O3 with 
concentrated caustic soda or potash, it goes into solution, but on 
dilution with considerable hot water Sb203 separates out again. On 
filtering this off, tetragonal crystals of NaSbOeg are deposited in the 
filtrate on cooling; they are, however, very unstable, and are decom- 
posed by standing in the air into sodium carbonate and antimony 
trioxide. By dissolving antimony trioxide in strong alkali, the ortho- 
antimonate is probably formed, 


Sb203+60H- — 2SbO3 +3H20, 


244 HYDROGEN SULFIDE GROUP 


which is hydrolyzed on dilution into metantimonite and alkali hy- 
droxide: 


SbO3=+H20 — 20H-+S8b0.7. 


The latter is decomposed by more water into trioxide and alkal: 
hydroxide; so that on adding to a solution of the trichloride either 
sodium hydroxide or carbonate, an almost quantitative precipita- 
tion of Sb2Oz will be obtained: 


2SbCl3 +6NaOH = 6NaCl+3H20+Sb203; 


2SbCl3+3Na2CO3 =6NaCl+3CO02+Sb203. 


Antimony oxychloride, SbOCI, contains the univalent group, SbOT, 
which is known as the antimony] group. 

Antimony oxychloride, therefore, can be regarded as antimonyl 
chloride. Antimonyl nitrate SbO(NOs), is also known, and anti- 
monyl sulfate, (SbO)2SO4. All these compounds are easily hydrolyzed 
into acid and oxide, so that they are rarely met with in the course of 
analysis, with the exception of antimony] chloride. ) 

The antimonyl compounds of certain organic acids (such as tar- 
taric acid) are very much more stable. : 

On boiling antimony trioxide with a solution of potassium acid 
tartrate, it goes readily into solution, forming the so-called ‘ tartar 
emetic,” 


2KHC4H406+Sb203 — H20+2K(SbO)C4H40¢, 


which is comparatively soluble in water. 100 cc. of water dissolve: 
5.3 gms. at 8.7°; 7.9 gms. at 21°; 12.2 gms. at 31°; 18.2 gms. at 50°, 
and 31.2 gms. at 75°. 
Not only antimonyl oxide, Sb2O3, but all antimonyl compounds, 
form a complex anion with tartaric acid or with a tartrate, thus, 
antimony] chloride dissolves in Rochelle salt, or in tartaric acid: 


C4H40¢6~+SbOC1 — Cl +[(SbO)CaH40¢)-; 
H2C4H406+SbOC1 — HCI+ H[(SbO)C4H 40g]. 


Tartar emetic, K(SbO)C,H.,0.-3H.O, is the most important antimony 
compound of commerce. Consequently it will be worth while to say a few 
words with regard to its behavior toward acids. 


ANTIMONY | 245 


If an aqueous solution of potassium antimony] tartrate is treated with 
hydrochloric acid, a white precipitate of antimony] chloride ,is formed, 


K(SbO)C,H,0.+2H* — H.C,H,0,+ K++ SbOCI, 
which readily dissolves in more hydrochloric acid, 
SbOCI+ 2HCl = H:0+S8bCl,, 


but, on the addition of more water, it is reprecipitated, etc. 

Sulfuric and nitric acids precipitate orthoantimonous acid from a solution 
of potassium antimony] tartrate; for the antimonyl compound, which is at 
first formed, is immediately decomposed by water, 


[(SbO)C,H,0.]-+Ht+2H.0 << H.C,H,O,+ H;SbO;. 


REACTIONS OF ANTIMONOUS COMPOUNDS IN THE WET WAY 


‘1. Water precipitates at first a basic salt which is changed into 
oxide by more water. 

2. Sodium Hydroxide, Ammonia, and Alkali Carbonates precipi- 
tate the amorphous hydrated oxide. 

3. Hydrogen Sulfide precipitates, from solutions which are not 
too acid, flocculent, orange-red antimony trisulfide: 


2Sb++*+3H2S — 6H*+SbeS3. 


As is indicated in the equation, the antimony trisulfide is in equilibrium with 
hydrogen ions; in 12-normal acid it is readily soluble (difference from arsenic). 
On the other hand, antimony sulfide is less soluble in acid than is tin sulfide 
and it can be precipitated from a solution which is normal in acid concentration. 
The trisulfides of both tin and antimony are somewhat more soluble than the 
pentasulfides. If a solution of antimony in concentrated hydrochloride is to 
be treated with hydrogen sulfide and the solution is first diluted, a precipitate 
of SbOCL is likely to form. This does no harm, as hydrogen sulfide will change 
the oxychloride into the less soluble trisulfide. If it is desired to precipitate 
the sulfide without any formation of oxychloride, the solution should be satu- 
rated with hydrogen sulfide before it is diluted, then diluted and again satu- 
rated with hydrogen sulfide. 


Antimony trisulfide is soluble in ammonium sulfide, forming a 
soluble thio salt: 
Sb2S3+3S7 — 2[SbS3]~. 


The triammonium salt has never been isolated, the mono salt, NH,SbS,, 
alone being known in the solid state. In solution, however, particularly in 
the presence of considerable ammonium sulfide, the SbS;= ion is probably 
present. 


246 HYDROGEN SULFIDE GROUP , 


If yellow ammonium sulfide is employed, ammonium salts of 
thioantimonic acid are obtained: 


SboS3+2(NH4)2S2 = (NH4)3SbS4-+ N H4SbSs3. 


If the solution of ammonium thioantimonite is boiled for a long time 
in the air, the red-colored oxysulfide is often precipitated: 


2(NH4)3SbS3+402 = 2(NH4)28203 +2NH3+H20+Sb2820. 


By boiling antimony chloride with sodium thiosulfate, the oxy- 
sulfide is also obtained, 


2Sb*t *+*++358203" > 4802 7 +Sb2820, 


which, on being warmed with ammonium sulfide, redissolves, forming 
the thio salt. | 

Antimony trisulfide is also soluble in caustic alkali, forming thio 
and oxythio salts: 


SbeS3+20H™ = H20+O8bS+S8S8bS-. 


These thio salts are decomposed by acids, precipitating antimony | 
trisulfide, with evolution of hydrogen sulfide, b 


2SbS3 +6H* — 3H2S Tf +Sb2Ss; 
2SbS4=-+6H* — 3H28 t +SbeSs; 
OSbS~+SSbS~+2H* — H20-+Sb2Ss. 


4. Zinc precipitates from solutions of antimony compounds metallic 
antimony. If a piece of platinum foil and a little zine are placed in 
an antimony solution containing hydrochloric acid, so that the two 
metals touch one another, the antimony is deposited on the platinum 
in the form of a black stain which does not disappear on removal of 
the zine (difference from tin). 

Lead will also cause the same reaction to take place (cf. p. 44). 

5. Potassium Iodide does not set free iodine when treated with 
an antimonous solution (difference from antimonic compounds). 


B. Antimonic Compounds 


Antimony pentoxide, Sb2O5, is formed as a yellow powder by oxi- 
dizing antimony with concentrated nitric acid and gently igniting 
the reaction product (antimonic acid). On strong ignition it loses 


oxygen and goes over into the very stable antimonous antimonate 
Sb204. 


ANTIMONY 247 


The pentoxide dissolves in concentrated hydrochloric acid, form- 
ing the pentachloride 


Sb20;+10HCl=5H20+2SbCls. 


If this solution is treated with water, a white precipitate of anti- 
monic oxychloride, SbO2Cl, is formed, which, by the addition of more 
water, is changed on warming into antimonic acid: 


SbCl; -+2H20 = 4HC1+Sb0.Cl, 
Sb0.Cl+2HOH @ HC1+H;Sb0u. 


Tartaric acid prevents the precipitation of the oxychloride, as 
with SbOCI (p. 244). Antimony pentoxide is an acid anhydride, 
and, like the corresponding P205, can be referred to three acids’ 
orthoantimonic acid, HsSbO4; metantimonic acid, HSbO3; and 
pyroantimonic acid, H4aSbeO7, which have all been isolated. The 
salts of the metantimonic and pyroantimonic acids are the most com- 
mon. ‘The trimetallic salts of the ortho acid have never been isolated, 
but the monometallic salts are known to exist. All antimonates, 
being salts of a weak acid, are very unstable, being easily hydrolyzed 
by water. ¢ 

If antimony pentoxide is fused with an excess of caustic potash, 
the product of the fusion probably contains the trimetallic salt of 
orthoantimonic acid. If, however, the melt is dissolved in a little 
water and allowed to crystallize, deliquescent crystals of potassium 
pyroantimonate, K4Sb207, are formed. 

The ortho salt, which is first formed, is decomposed by water as 
follows: 

2K3SbO4+ H20 @ 2KOH+ KaSbe207. 


By the action of considerable cold water (or more quickly by 
rapid boiling with less water) this deliquescent salt is gradually 
changed into the acid salt, losing KOH, 


K4Sbe07+2H20 @ 2KOH+Ke2H2Sb20;7, 


which separates out with 6 molecules of water as a granular powder, 
difficultly soluble in cold water. It dissolves to a considerable extent 
in water at about 40°—50° C.; this solution is used in testing for sodium, 
as the sodium salt is very much more insoluble in water. 

On boiling the granular potassium salt for a long time with con- 
siderable water, it gradually takes on water, forming the easily soluble 
monometallic salt of orthoantimonic acid, 


K2H2Sb207+H20 @ 2KH2SbO04, 


2A8 HYDROGEN SULFIDE GROUP 


which is obtained, on evaporating the solution, as a gummy mass 
of the composition 2KH2SbO4+H20, but on boiling the aqueous 
solution for a long time, more KOH is lost, with the formation of 
amorphous orthoantimonic acid: | 


KH2SbO4+ H20 @ KOH+H3SbO.. 


All antimonates are decomposed by acids, amorphous antimoniec 
acid separating out. 

The gummy, monometallic salts give an amorphous precipita- 
tion with sodium salts, gradually becoming crystalline, while the 
potassium pyroantimonate gives a crystalline precipitate immediately. 

The relations of antimonous to antimonic salts and of antimonites to anti- 
monates are, on the whole, similar to the corresponding relation with arsenic. 
In the higher state of oxidation the acid properties are more pronounced than 


the basic properties, but antimony, being lower in same group of the periodic 
table, is more basic than arsenic. 


REACTIONS OF ANTIMONIC COMPOUNDS IN THE WET WAY 


1. Hydrogen Sulfide precipitates from fairly acid ‘solutions the 


orange-red pentasulfide: 
+44 


28b 4 +5HeS a 10Ht+SboSs. 


Antimony pentasulfide is soluble in 12-normal hydrochloric acid, 
forming antimony trichloride, with deposition of sulfur and evolution 
of hydrogen sulfide: 


Sb2S5;+6HCl = 3H2eS+S8e2 +28bCls. 
It also dissolves (like the trisulfide) in alkali sulfides, and in alka- 


lies, but not in ammonium carbonate. By treatment with an alkali 
sulfide, a thio salt is obtained, 


Sb2S5+38~ — 28bS.4°, 


which is decomposed by the addition of acids, forming the insoluble 
pentasulfide with evolution of hydrogen sulfide: 


2SbS4-+6H* — 3H2S fT +Sb2Ss. 
Alkalies dissolve the pentasulfide, forming thio and oxy-thio salts: 
Sb2S5+60H~ — Sb84=+8-SbO3 +3H20. 


2. Hydriodic Acid reduces antimonic compounds in acid solu- 
tions, with separation of iodine (difference from antimonous com- 
pounds) : 

Shi ttt t+91- + Spttt+tp. 


a a | a 


ae og 
Pd! 


TIN 249 


_ 3. Nascent Hydrogen.—By treating any solution which ¢éontains 
antimony with nascent hydrogen, stibine is formed: 


Sb203+6H2 =3H20+2SbHs f , 
Sb205-+8H2 = 5H20+2SbHs } . 


If the stibine is generated in a Marsh apparatus (cf. p. 235), and the gas 
is conducted through a red-hot glass tube, a mirror of metallic antimony will 
be deposited, as with arsenic. But as stibine is much more unstable than 
arsine, and the antimony itself is much less volatile, the mirror is found nearer 
the heated place than is the case with arsenic—sometimes before the hottest 
part of the tube is reached—as the decomposition of the stibine takes place 
at a much lower temperature than with arsine. 

If the stibine is allowed to escape from the tube with the hydrogen, it 
' burns with a pale greenish-white flame to water and antimony trioxide. If 
a piece of glazed porcelain is held directly over the flame, a deposition of 
metallic antimony is obtained which is unaffected by a solution of sodium 
hypochlorite (difference from arsenic). 

If stibine is allowed to act upon a solution of silver nitrate, a black pre- 
cipitation of silver antimonide is thrown down: 


SbH;+3Agt =Ag;Sb+3H". 


Solid silver nitrate is turned yellow at first, then black; exactly the same 
as by arsine (cf. p. 238). 


REACTIONS OF ANTIMONY IN THE DRY WAY 


Antimony compounds impart to the flame a pale, greenish-white 
color. Heated with sodium carbonate on charcoal, a brittle metal- 
lic button is obtained, surrounded by a white incrustation. 

Compounds containing oxygen are reduced in the upper reducing 
flame to metal, which is volatile and burns in the upper oxidizing 
flame to trioxide; the latter can be deposited on a glazed porcelain 
surface. If the deposit is moistened with silver nitrate solution, 
and ammonia blown upon a it becomes black, owing to the separation 
of metallic silver: 


Sb203-+4AgNO03+4NH3+2H20 = 4NH4NO3 +Sbe20; +4Ag. 


TIN, Sn. At. Wt. 118.7 
' Sp. Gr. =7.29. M. Pt. =232°C. B. Pt.=about 1500° C. 


Occurrence.—Tin does not occur free in nature, but mostly in 
‘the form of the dioxide, as tetragonal tinstone, or cassiterite, isomor- 
phous with rutile (TiOg), zircon, and polianite (MnOz). 

Tin is a silver-white metal, which is ductile and malleable at ordi- 


250 HYDROGEN SULFIDE GROUP ~ 


nary temperatures, but at low temperatures and near the melting- 
point it is so brittle that it can be powdered. In order to pulverize 
tin, heat it in a porcelain dish till it melts, remove- the flame and 
quickly crush the substance with a pestle. It soon cools to about 
200°, becomes brittle and yields a fine powder. 
Tin is soluble in hot, concentrated hydrochloric acid with evolu- 
tion of hydrogen: 
Sn+2H*t — Sntt+He. 


In the presence of platinum the solution takes place more quickly 


and at a lower temperature. Dilute hydrochloric acid dissolves tin, - 


but very slowly. 
Nitric acid, of sp.gr. 1.2 to 1.3, does not dissolve tin, but oxidizes 
it to metastannic acid: 


3Sn+4HNO3+H20 =3H2Sn03+4NO f. 


Cold, dilute nitric acid dissolves the metal very slowly, without any 


evolution of gas, forming ammonium and stannous nitrates. 


In this reaction the tin is given two positive charges and a part of the 
nitrogen of nitric acid is reduced from a positive valence of five (toward oxy- 
gen) to a negative valence of three (toward hydrogen) thereby losing eight 
charges. Thus one molecule of nitric acid oxidizes four atoms of tin and 
eight more molecules of nitric acid are required to form stannous nitrate and 
one to form ammonium nitrate. The whole reaction may be expressed as 
follows: 


4Sn-+ 10HNO3 = 4Sn(NOs) 2 +NH4NO3 +3H20. 
Aqua regia dissolves tin, forming stannic chloride: 
3Sn+4HNO3+12HCl1=4NO 7 +8H20+3S8nCh. 


Tin dissolves in dilute sulfuric acid very slowly, but readily in 
hot concentrated acid, forming stannic sulfate, with evolution of 
sulfur dioxide: | 


Sn+4H2SO04= 2802 fT +4H20+Sn(S04)2. 


Tin forms two oxides: stannous oxide, SnO; and stannic oxide, 
SnOz. Salts are known corresponding to both these oxides—stannous 
and stannic salts. The former contain bivalent tin and the latter quad- 


rivalent tin. Tin is more basic in its properties than antimony, — 


but stannous hydroxide is slightly amphoteric (cf. p. 52) and the 
acid properties of stannic hydroxide are more pronounced. 


| —— _ 


TIN 251 


Stannous Compounds 


Stannous oxide (according to the way it is prepared) is either an 
olive-green or a black powder, which, on being warmed in the air, 
like all stannous compounds readily changes to stannic oxide. By 
dissolving stannous oxide (or, better still, the metal itself) in hydro- 
chloric acid, stannous chloride is obtained, which is the most im- 
portant of all the stannous salts. This salt, with two molecules of 
water of crystallization, SnCle+2H20, is the so-called “ tin salt” of 
commerce. 

Fresh crystals of ‘ tin salt ”’ will dissolve clear in a little water; 
if more water is added the solution becomes turbid, owing to the 
formation of a basic salt, 


SnClo+H20 2 HCl+Sn(OH)CI, 


- 


which is readily soluble in hydrochloric acid. 
The clear concentrated solution also becomes turbid on stand- 
ing in the air, owing to the formation of the same basic salt. 


28nClo-+0-+H20 > 2Sn(OH)CI1+Cle, 


The chlorine, however, is not set free, but unites with some of the 
stannous ions, forming stannic ions: 


Sn**--Clh —s Sat t?+-LOCr-. 


If tin tetrachloride is treated with metallic tin, the latter goes 
into solution and the former is reduced to stannous chloride: 


SnCl4+Sn = 28nCle. 


Consequently, in order to keep a solution of stannous chloride 
in the stannous condition, hydrochloric acid should be added to pre- 
vent the formation of the basic salt, and metallic tin to keep the solu- 
tion reduced. 

Such a solution constantly grows more concentrated, owing to 
the gradual solution of the tin. In order to keep a solution of stannous 
chloride at a definite concentration (only necessary for purposes of 
quantitative analysis) the hydrochloric acid solution should be kept 
out of contact with air in an atmosphere of carbon dioxide without 
the addition of metallic tin. | 

Nearly all stannous compounds are colorless; the oxide (as already 
mentioned) is black and the sulfide dark brown. 


252 HYDROGEN SULFIDE GROUP 


REACTIONS IN THE WET WAY 


1. Potassium and Sodium Hydroxides produce a white precipi- 
tate of gelatinous stannous hydroxide, 


Snt+++20H- > Sn(OH)o, 


which is readily soluble in an excess of the precipitant, forming alkali 
stannite: | 


Sn(OH)2+OH~ — H20+H8n0-2_, 


The hydroxide is also readily soluble in hydrochloric acid. 

The alkaline solution of an alkali stannate is often changed to 
brownish black or black (particularly on warming, or when very con- 
centrated caustic alkali is used), owing to the separation of either 
metallic tin or stannous oxide (ef. p. 170): 


From dilute potassium hydroxide solutions there gradually sepa- . 


rates on standing, or more rapidly on heating, the black monoxide, 
HSnO2 — OH +Sn0O; 


and from quite concentrated alkali the precipitate is almost wholly 
black metallic tin: 
2HSnO2 —SnO3 +H20+Sn. 


2. Ammonia and Alkali Carbonates precipitate the white hydrox-. 
ide, which is not absolutely insoluble in an excess of the precipitant: — 


Sn**+20H- — Sn(OH)>; 
Snt*+C0O37+H20 — COe 7 +$n(OH)o. 


A large quantity of tartaric acid, more than is required in the 
case of antimony, prevents the precipitation. 

3. Hydrogen Sulfide produces (in solutions which are not too 
acid) a brown precipitate of stannous sulfide, 


Snt++H.S > 2H*+SnS, 


readily soluble in strong hydrochloric acid; therefore no stannous 
sulfide is precipitated if the solution is very acid. After diluting 
a strongly acid solution with water, however, stannous sulfide is 
completely precipitated on saturating the solution with hydrogen 


sulfide gas. 
Stannous sulfide is insoluble in ammonia and ammonium carbon- 


ate (difference from arsenic); also in colorless ammonium sulfide (dif- 





oe 


. * f 
~~). os ae 


TIN 253 


ference from arsenic and antimony); but is readily soluble in yellow 
ammonium sulfide, forming ammonium thiostannate: 


SnS+ (NH4)2Se = (NH4)2SnSs3. 


If the solution of ammonium sulfostannate is acidified with any 
acid, yellow stannic sulfide is precipitated: 


SnS3 +2H ex HS T +SnSo. 


. 4. Mercuric Chloride produces in solutions of stannous salts a 
white precipitate of mercurous chloride: 


2HgCle+Sn* *t > Snt*+++2Cl-+H¢g.Clo. 


But if the stannous chloride is present in excess, the mercurous 
chloride will be reduced to gray mercury: 


Hg2Cle+Sn** — Sn*t*+++2Cl-+2Hg¢g. 


5. The Gold Test is much more sensitive. If to a solution of 
gold chloride a solution containing a trace of stannous chloride is added, 
finely divided metallic gold will be precipitated, 


2Aut *++38ntt — 3Sntt+t+42Au, 


which appears brown by transmitted light, and bluish-green by reflected 
light. 

6. Metallic Zinc precipitates tin from both stannous and stannic 
solutions as a spongy mass, which adheres to the zinc: 


Sn? ?-Zn = Zn t-+-Sn: 


The finely-divided, spongy metal is easily soluble in strong hydrochloric 
acid; the experiment must not, therefore, be made in strongly acid solutions, 
as the tin is loosened from the zinc by the violent evolution of hydrogen and is 
redissolved by the acid. The test is best made by adding a drop of the (not 
too acid) solution to a piece of platinum foil, and then placing a piece of bright 
zinc so that it comes in contact with both the golution and the platinum. 
The tin is precipitated partly on the zinc and partly on the platinum,* in 
the form of a gray stain, which disappears from the latter as soon as the zine 
is removed, provided the solution is still acid (difference from antimony), 
If the zinc is kept in contact with the acid until the evolution of hydrogen ceases, 
the tin stain will not disappear from the platinum, because all the acid has been 
used up. On adding a few drops of concentrated hydrochloric acid to the 
platinum, the stain quickly disappears with an evolution of hydrogen. The 
reason why the tin is deposited on the platinum notwithstanding the pres- 





*In weakly acid solutions tin is precipitated chiefly on the zinc; in strongly 
acid solutions, chiefly on the platinum. 


254 - HYDROGEN SULFIDE GROUP 


ence of acid is that a galvanic current is formed by the contact of the zine 
with the platinum, which flows from the zine to the platinum; the platinum 
thus serves as a cathode, and the tin is deposited upon it. On removing the 
zinc the current stops and the stain disappears. 


7. Metallic Lead, as is evident from its position in the electro- 
motive series (cf. p. 41), will reduce tin to the metallic state, but the 
reaction will stop as soon as the concentration of Pb** becomes nearly 
as large as that of Sn**; when the concentration of Pb** is greater 
than that of Sn** the reaction will take place in the reverse direction. 
It requires a relatively low concentration of H* to stop both of these 
reactions, as both tin and lead are higher in the series than hydrogen. 
In acid solution, however, metallic lead will reduce Sn***+* to Snt* 
(cf. p. 44). 


Stannic Compounds 


The stannic compounds (which are all colorless, with the excep- 
tion of the yellow sulfide, SnSz2) cannot be obtained by the solution 
of the oxide, SnOz, from which they are derived, because the oxide 
is attacked with difficulty by acids. They are obtained indirectly 
from metallic tin or from stannous compounds. 

The simple stannic compounds are all, more or less readily, com- 
pletely hydrolyzed by water, so that the analyst almost never meets 
with them. The nitrate, Sn(NOs)4, and the sulfate, Sn(SO4)o, are 
quickly decomposed in the cold, into acid and stannic hydroxide. 
The halogen compounds are more stable, and are decomposed only 
by boiling with considerable water. For the following reactions, 
therefore, we will assume that we have a solution of stannic chloride 
to work with. | 

Stannic chloride is a colorless liquid, which fumes in the air and 
boils at 120° C. On adding a little water it solidifies, forming crystals 
of monoclinic hydrates, SnCl4a+3He2O, SnCl4a+5H20, SnCl4a+8H20, 
of which the salt with 5H2O is used commercially as a mordant in 
dyeing. ; 

On adding more water to these hydrates they dissolve, forming 
a clear solution, which on boiling (the freshly-prepared, dilute solution) 
gradually becomes turbid, owing to the precipitation of voluminous 
stannic hydroxide: 


SnCl,+4HOH = 4HCl+Sn(OH)4q. 
If the solution is very dilute it becomes turbid in the cold. The 


stannic acid thus formed is not precipitated quantitatively, either 
in the cold or on boiling, because a tonsiderable amount remains in 


TIN 255 


the hydrosole form. By ‘salting out” the hot solution (best with 
ammonium nitrate), the stannic acid may be completely precipitated. 
A solution of stannic chloride can be most readily obtained for 
analytical purposes by chlorinating or brominating a solution of 
stannous chloride. 
On ‘adding chlorine to a solution of stannous chloride, stannic 
chloride is formed in the cold: 


SnCls 4. Cle = SnCl.. 


As, however, chlorine is colorless in a dilute solution, it is difficult 
to tell when the oxidation is complete; it is more easily ascertained 
if bromine is used. 

On adding strong bromine water, drop by drop with constant 
stirring, to a solution of stannous chloride, the brown color will dis- 
appear as long as any stannous chloride remains unchanged, and the 
solution becomes colored by the bromine only when the oxidation is 
complete. ‘The solution then contains a mixture of stannic chloride 
and stannic bromide: 


2SnClo+2Bre=SnCl+SnBra. 


Just as platinum tetrachloride combines with hydrochloric acid 
to form chloroplatinic acid, so tin tetrachloride unites with hydro- 
chloric acid, forming chlorostannic acid, HeSnCle, and yields, like 
the former, beautifully crystalline, easily soluble salts with the alka- 
lies, of which the ammonium salt (NH4)2[SnCl¢] is an article of com- 
merce, being known as “ pink salt.’ The above-mentioned stannic 
chloride, SnCl4, is sometimes designated as a-stannic chloride, to 
distinguish it from a compound (which we shall soon study) known 
as 6-stannic chloride (stannyl chloride). 


Reactions of the a-Stannic Compounds 


1. Hydrochloric and Sulfuric Acids produce in moderately con- 
centrated solutions of stannic chloride no precipitation, even on long 
standing (difference from #-stannic compounds). In very dilute 
sulfuric acid solutions a precipitate of basic sulfate is sometimes 
obtained. In very dilute hydrochloric acid solutions, also, a slight 
turbidity is often formed, which increases on boiling the solution: 


SnCly+4HOH = 4HCI+Sn(OH)s. 


2. Potassium and Sodium Sulfates produce no precipitation in 
the cold (difference from stannyl chloride), but on boiling all the tin 
is precipitated as hydroxide. 


256 ~ HYDROGEN SULFIDE GROUP 


3. Potassium or Sodium Hydroxide.—On adding caustic alkali 
to a solution of a stannic salt, a voluminous, gelatinous, white pre- 
cipitate is obtained: 


SnCli +40H- — 4CIl--+Sn(OH),. 


The precipitate has the above formula when dried in the air, and 
the formula He2SnOs if dried over sulfuric acid. 

The precipitate dissolves readily in an excess of alkali hydroxide, 
forming salts which are not derived from either of the above com- 
pounds, but from He-[Sn(OH).6|, which has itself never been isolated: 


(SnOH),+2KOH = K2[Sn(OH)<]. 


The hydroxide dissolves in ammonia also, but only in the absence 
of ammonia salts. 

By dissolving in alkali, stannic hydroxide behaves as an acid, 
and according to Bellucci and Parravano,* the hexahydroxystannic 
acid stands in the same relation to chlorostannic acid as hexahydroxy- 
platinic acid to chloroplatinic acid: 


Ho2[PtCle] Ho2[SnCle] 
H.[Pt(OH)c] H.2[Sn(OH)g]. 


The salts of hexahydroxystannic acid are designated briefly as 
stannates, or a-stannates, to distinguish them from the 6-stannates 
or metastannates, which are derived from the polymer (H2SnQOs3)5 
(see p. 257). 

The ready solubility of e-stannic acid in cold dilute mineral acids 
is very characteristic. It dissolves promptly in hydrochloric, nitric, 
and sulfuric acids, behaving, in this respect, as a base. By boiling the 
dilute acid solution (particularly the sulfuric acid solution) stannic 
acid is reprecipitated, which is soluble in cold dilute acids provided 
the boiling has not been continued too long. In the latter case the 
6-stannic acid is formed, which is insoluble in dilute acids. 

4. Potassium Carbonate precipitates stannic acid from stannic 
chloride solutions; the precipitate is completely soluble in an excess 
of the reagent (difference from 6-stannic acid): 


SnCly -- 2K2CO3 a 2H2O =4KCl ao 2CO2 ad Sn(OH),. 


5. Sodium Carbonate behaves similarly, but the precipitate is 
not so easily soluble in an excess. 
6. Ammonia precipitates stannic acid from a solution of stannic 


* Z. anorg. Chem., 45 (1905), p. 156. 





TIN | 257 


chloride; tartaric acid prevents the precipitation (difference from 
B-stannic acid). 


Reactions of the 6-Stannic Compounds 
(Metastannates) 


By the oxidation of metallic tin with hot nitric acid of sp.gr. 1.3 
stannic nitrate is first formed, which, by boiling with water, is com- 
pletely hydrolyzed, forming nitric acid and metastannic acid. 

Metastannic acid is a white powder insoluble in nitric acid, and 
when dried over sulfuric acid has the formula H2SnO3. This is of the 
same empirical composition as the hydroxide precipitated by treating 
a stannic chloride solution with alkalies, though differing essentially 
from it in many reactions. 

While the a-stannic acid (as already mentioned) is easily soluble . 
in dilute mineral acids, the 6-stannic acid is almost insoluble therein. 

1. If the B-stannic acid is treated for a short time with concen- 
trated hydrochloric acid, a chloride is formed which a insoluble in 
hydrochloric acid, but readily soluble in water. The solution con- 
tains the so-called 8-stannic chloride (though the designation stannyl 
chloride would be more suitable) of the composition, Sn5O5Cle(OH)s.* 

2. On treating the aqueous solution of stannyl chloride with 
hydrochloric acid, almost all the tin is reprecipitated in the form of 
a highly chlorinated compound of the composition Sn505ClufOH)6+ 
4H20.t , 

3. If an aqueous solution of stannyl chloride is heated to boiling, 
almost all the tin is precipitated as 6-stannic acid, which is insoluble 
in dilute acids. 

This differing behavior of the two acids, as well as of the two 
chlorides, can be explained as follows: Silicic acid, which is closely 
related to stannic acid, exists in innumerable silicates.in different 
polymeri: forms. Thus, with the minerals of the pyroxene and 
amphibole groups, wollastonite, CaSiOs, is a derivative of ordinary 
meta silicic acid and tremolite, CaMgs3Si4Oi2, is a derivative of 
' (He2SiO3)4.f It is highly probable that the stannic acid can exist 
in analogous polymers. One of these polymers apparently corresponds 
to the composition (H2SnQs3)s. 

If such a compound is treated with hydrochloric acid, the hydroxyl 





*R. ENGEL, Chem. Zig., 1897, pp. 309 and 859. 
+ Weser, Jahresber., 1869, 244, and Pogg. Ann., 122, 358. 
t Grorn, Tabellarische Uebersicht d. Min., 1898, p. 148. 


258 HYDROGEN SULFIDE GROUP 


groups will, first, be replaced by chlorine, and a compound will be 
obtained containing tin, oxygen, and chlorine, e.g., Sns5O5Clro. 

This hypothetical compound of the 6-stannic acid is hydrolyzed, 
forming different chlorides of varying solubilities. Thus R. Engel 
found that the chloride SnsOs5Cle(OH)g is soluble in water; and 
Weber showed that from an aqueous solution of the latter, hydro- 
chloric acid precipitates the compound Sn5O5Cl4(OH)6+4H20. 

The reaction which takes place on dissolving the §-stannie acid 
in hydrochloric acid and water may be expressed satisfactorily by 
the following equations: 


Sn;05(OH)10+10HCl=10H20+Sn;05Clio (insoluble in HCl); 


8-stannic acid 


or 
Metastannic acid 


S$n505Clio+8H20 — 8HCI+S8n505Cle(OH)s (soluble in water) ; 
Sn505Cle(OH)s+2HCl @ 2H20+S$n505Cl4(OH)¢ (insoluble in HCl). 


On boiling the aqueous solution, complete hydrolysis takes place: 
Sn505Clo(OH)s+2HOH =2HCI+Sn;505(OH) 0. 


If the 6-stannic acid is treated for a long time with concentrated 
hydrochloric acid, the Sn5Os5 group is finally broken down, and the 
tin goes into solution in the form of ordinary a-stannic chloride: 


$n50;5(OH)19+20HCI = 15H20+5SnCly. 


Further reactions of stannyl chloride (6-stannic chloride) are: 

4. Sulfuric Acid precipitates from solutions of stannyl chloride 
white stannyl sulfate, which on being washed with water is completely 
changed to $-stannic acid (difference from a-stannic chloride). 

5. Potassium and Sodium Sulfates cause the same reaction as 
sulfuric acid. 

6. Potassium Hydroxide throws down in solutions of stannyl 
chloride a voluminous precipitate of 6-stannic acid, which does not 
dissolve in an excess of the concentrated precipitant, but forms a 
6-stannate easily soluble in dilute caustic potash solution: 


(a) Sns05Clo(OH)s+2KOH =2KCl+Sn;0;(OH) 10; 


B-stannic acid 


(b) Sn505(OH) 10 +2KOH = 2H20+8n505(OK)2(OH)s. 


potassium #-stannate 


By long treatment of the potassium #-stannate with concentrated 
caustic potash, it gradually goes into solution, forming a-potassium 


TIN 259 


stannate. This change takes place more readily by fusing 6-stannic 
acid with solid potassium hydroxide in a silver crucible. 

If a dilute solution of a mineral acid is added to the 6-potassium 
stannate, a voluminous precipitate is formed, consisting partly of 
B-stannie acid (insoluble in an excess of mineral acids) and partly 
of a-stannic acid (readily soluble in an excess of the acid). ‘The latter 
compound is formed when a very concentrated solution of caustic 
potash was used in forming the potassium salt. 

7. Ammonia also precipitates 6-stannic acid, even in the pres- 
ence of tartaric acid (difference from a-stannic chloride). 

As we have seen, the a-compounds may be readily changed into 
B-compounds and conversely. The dilute aqueous solutions of the 
a-compounds are gradually changed, at the ordinary temperature, into 
B-compounds, but more quickly on boiling; thus stannic chloride changes 
to stannyl chloride: 


58nCl4+13H20 =18HC1+Sns505Cle(OH)s. 


The B-compounds are changed into a-compounds by boiling with 
oncentrated hydrochloric acid or with concentrated caustic potash. 


8. Hydrogen Sulfide precipitates (from not too acid solutions) 
yellow stannic sulfide from both the.a- and the 6-compounds: 


SnCl,+2Has a= 4HCl+- SnS2; 
Sn505Cl2(OH)g+10H2S @ 2HCl1+13H20+5SnSo. 


From £-stannic solutions hydrogen sulfide produces a precipitate, but 
very slowly, the SnS, remaining largely in the hydrosole form. By the addition 
of salts it is coagulated, and separates out in a flocculent form, usually mixed 
with g-stannic acid. If the stannyl chloride solution is heated on the water- 
bath in a pressure flask, the tin is quickly precipitated as greenish-yel'ow 
sulfide. 


Stannic sulfide is soluble in hydrochloric acid; hydrogen sulfide 
will cause no precipitation, therefore, if the solution is very acid. If 
such a solution, saturated with hydrogen sulfide, is largely diluted, 
the sulfide will precipitate out. 

The yellow sulfide is the thio-anhydride of the thio-acid; it dis- 
solves, therefore, in alkali sulfide, forming salts soluble in water: 


SnS2+ (NH4) 28 — (NH4)2Sn8sz. 
Acids precipitate from such a solution the yellow stannic sulfide 


SnS37-+2H* — HeS+SnSo. 


260 HYDROGEN SULFIDE GROUP 


The sulfide is insoluble in ammonia and ammonium carbonate 
(difference from arsenic). By means of concentrated nitric acid it 
is easily oxidized to 8-stannic acid; or by roasting in the air it can 
be completely changed to tin dioxide. 

The sulfide obtained in the dry way, known as “ mosaic gold,” 
is not attacked by nitric acid, and is also insoluble in alkali sulfides. 
It dissolves on treatment with aqua regia, forming stannic chloride 
with separation of sulfur. It is most readily brought into solution 
by fusing with sodium carbonate and sulfur (see below). 

9. Mercuric Chloride produces no precipitation in solutions of — 
stannic salts. 


Tin dioxide as it occurs in nature, and the artificially produced oxide 
after strong ignition, are both insoluble in all acids. They can be brought 
into solution by the following methods: ! 

1. Fusion with sodium carbonate and sulfur; 
2. Fusion with caustic potash or soda; 

3. Fusion with potassium cyanide; 

4. Reduction by hydrogen at a high heat. 

(1) Fusion with Sodium Carbonate and Sulfur.—Place the dry substance 
in a small porcelain crucible, mix with six times as much calcined sodium 
carbonate and sulfur (equal parts mixed together), cover the crucible, and heat 
over a small flame until the excess of sulfur has distilled off and burned. This 
operation requires about twenty minutes. Allow the crucible to cool and 
then treat its contents with hot water, and filter if necessary: . 


2S8n0, +2Na.CO; +95 =3S0, T +2Na.Sn8;+2CO, tT ° 


If iron, lead, copper, or any other metal that forms a sulfide insoluble in 
water and ammonium polysulfide is present, it remains undissolved as sulfide, 
and is separated from the tin by filtration. 

(2) Fusion with Sodium Hydroxide.—Melt the sodium hydroxide in a silver 
crucible, placed within a porcelain crucible to protect the silver from the 
injurious action of the gas flame. When the fusion has become quiet, owing 
to the expulsion of all the excess water, cool somewhat and add the finely- 
powdered, dry tin dioxide. Heat until the fusion is clear. After cooling 
dissolve the melt in water: 


SnO, +- 2Na0H = NaSnQ; + H.0. 


Stannic oxide is not completely attacked by fusion with sodium or potassium 
carbonate. 

(3) Fusion with Potassium Cyanide-—Melt some potassium cyanide in a 
porcelain crucible, add the powdered stannic oxide and fuse the mixture until 
‘the separated tin has melted together: 


Sn0.+2KCN =2KCNO-+Sn. 


After cooling, extract the mass with water, filter off the tin, flatten it into 
foil and dissolve it in concentrated hydrochloric acid. 
(4) Reduction in a Slream of Hydrogen.—Place the substance in a platinum 


- 


‘9 


a ae 


GOLD 261 


boat and insert the latter in a glass tube, open at both ends, which is made 
of difficultly fusible glass. Pass hydrogen through the tube in the cold until 
the air has been driven out, and then heat to dull redness until no more water 


is formed: 
SnO,.+4H =2H.0+S8n. 


Finally cool and dissolve the metal in hydrochloric acid. 


REACTIONS IN THE DRY WAY 


If a tin salt is heated with soda (or potassium cyanide) on char- 
coal, usually only a small malleable button is obtained, which, on 
taking away the flame, is immediately covered with a white coating 
of oxide. This can be observed when the flame is allowed to play 
upon the fusion. If the product is crushed in an agate mortar, a 
small flake of metallic tin is obtained, which can be distinguished 
from silver and lead by its insolubility in nitric acid, and by its solu- 
bility in hydrochloric acid. This reaction is particularly suitable 
for the charcoal-stick test. The borax bead which has been colored 
pale blue by copper becomes a transparent ruby red in the reducing 
flame if a trace of tin is added. This is a very sensitive reaction. 


GOLD, Au. At. Wt. 197.2 
Sp. Gr.=19.33. M. Pt. =1063° C. 


Occurrence.—Gold usually occurs native in quartz and in river 
sands; also as telluride of gold in sylvanite, (AuAg)2Tes, and in 
- nagyagite, (PbAu)2(TeSSb)3, and is found in small amounts in many 
pyrite and other sulfide ores. 

Metallic gold is of a yellow color and melts without being oxidized. 
It is the most ductile of all metals, and may be hammered into 
exceedingly thin leaves, which are transparent, with a bluish-green 
color. 

Commercial gold is usually alloyed with copper, silver, or with 
both metals.. Pure gold is designated as 24-carat gold or !9°/j000 
fine. Fourteen-carat gold contains 14 parts of gold to 10 parts of 
other metal and 18-carat gold contains 18 parts of gold and 6 parts 
of alloy; the former, therefore, contains 58.3 per cent and the latter 
75 per cent of pure gold. 

The proper solvent for gold is aqua regia, but it is also soluble in 
bromine and chlorine water, forming a trihalogen compound: 


2Au+2HNO3+6HCl=4H20+2NO T +2AuCls, 
Au+3Br=AuBrs. 


262 HYDROGEN SULFIDE GROUP 


Gold is not» attacked by minerals. It forms two oxides: aurous 
oxide, Au2O; and auric oxide, AugQ3. 

Both of these are exceedingly unstable; on gentle ignition they 
lose oxygen and are changed to metal (a property common to all 
“noble ”’ metals). 

All gold salts are unstable; even the most stable salt of all, the 
chloride, AuCls, is changed by gentle ignition into yellowish-white 
aurous chloride, AuCl: 

AuCls = AuCl +Cle. 


On stronger ignition the last atom of chlorine is lost, and the yellow 
metal itself is left behind. 

Aurous chloride, AuCl, is insoluble in water, but on being boiled 
with water for some time, or very gradually in the cold, it is changed 
to auric chloride, with deposition of metal: 


3AuCl= AuCl3+2Au. 


The solution obtained by dissolving gold in aqua regia always 
contains auric chloride, so that only the reactions of aurie com- 
pounds are of interest to the analytical chemist. Auric chloride 
unites with hydrochloric acid, formiug chlorauric acid, 


AuCl3-+HCl=H[AuCly], 


which yields beautifully crystalline salts. 

Gold chloride is soluble in ether and can be extracted from its 
aqueous solutions by means of this solvent. : 

Auric salts are mostly yellow and readily soluble in water. The 
sulfide is black and soluble only in aqua regia. 


REACTIONS IN THE WET WAY 


A solution of chloraurie acid, H{[AuCl4], should be used for these 
reactions. 

1. Potassium or Sodium Hydroxide.—If caustic alkali is cau- 
tiously added to a concentrated gold solution, a reddish-brown, volu- 
minous precipitate of auric hydroxide is obtained, which looks exactly 
like ferric hydroxide. If more caustic alkali is added, however, the 
gold hydroxide redissolves, forming alkali aurate: 


Autt+*++30H-— Au(OH)3; Au(OH)3+OH~ — 2H20+[AuO,]-. 


If the bright-yellow solution of potassium aurate is carefully acidi- 
fied with nitric acid, a precipitate of reddish-brown auric acid is thrown 


GOLD 263 


down, which is soluble in nitric acid, but is reprecipitated, for the 
most part, by diluting and boiling. 

As a rule, potassium hydroxide yields no precipitate in solutions 
of gold chloride, because the gold solution is usually so dilute that 
the amount of alkali added is sufficient to form the aurate at once. 

2. Ammonia throws down a yellowish mixture of gold imino- 
chloride, Au(NH)Cl, and gold iminoamide, Au(NH)NHa2, 


AuCl3+3NH3 — 2NH4Cl+Au(NH)Cl1, 
AuCl3-+5NH3 — 3NH4Cl+Au(NH)NHa2, 


which explodes in a dry condition on warming or by concussion (ful- 
minating gold.*) 

The most important reactions for the detection of gold are those which 
depend upon the extreme readiness with which the auric compounds are 
reduced. Auric compounds are strong oxidizing agents. 

3. Ferrous Salts precipitate at ordinary temperatures from neutral 
or acid solutions all the gold as a brown powder (difference from 
platinum): : 
Aut+++43Fet+— 3Fet+++Au. 


4. Oxalic Acid precipitates all of the gold in the cold, but more 
quickly on warming (difference from platinum): 


2Aut 7 +41 3H C204 nex. 6Ht+6CO2 T +2Au. 


The presence of considerable acid prevents this reaction. 
5. Arsine and Stibine precipitate gold completely: 


2Aut*++AsH3+3H20 — H3AsO3+6H*+2Au. 
6. Sulfurous Acid reduces gold solutions: 
2Aut**+3S802+6H20 — 12H++3S804- +2Au. 
7. Stannous Chloride causes the following reaction to take place:T 


2Autt+++3Snt+— 3Snt*t*+2Au. 


If the solution tested is strongly acid with hydrochloric acid, the precipi- 
tate is pure gold and has the characteristic dark-brown color of the finely 
divided metal. In very dilute weakly acid solutions the so-called purple 
of Cassius is thrown down, which consists of colloidal gold and tin hydroxide.f 

Purple of Cassius is soluble in ammonia and in dilute caustic potash solu- 





* Rascuia, Ann. Chem. Pharm. 235 (1886), 325. 
+ THEopor Doérina, Chem. Centralbl., 1900, I, p. 735. 
t Zstamonpy, Ann. Chem. Pharm., 301 (1898) 361. 


264 HYDROGEN SULFIDE GROUP 


tion, forming reddish liquids. These solutions when cold remain clear for a 
long time and can even be boiled wlthout decomposition. As the solution 
is concentrated a flocculent precipitate is formed which will dissolve on the 
addition of more ammonia. 

The brown coloration can be distinctly seen if 0.3 mg. gold is dissolved 


in 100 cc. of the solution; if less than 0.3 mg. of gold is present, only a yellow-— 


ish coloration is obtained. 


8. Hydrogen Peroxide* in alkaline solution immediately pre- 
cipitates the gold as finely divided metal: 


2Au***+3H202+60H — 6H20+302 fT +2Au. 


The precipitated metal appears brownish-black by reflected light, 
but bluish-green by transmitted light; 0.03 mg. gold in 10 ce. of liquid 
suffices to give a reddish coloration with a bluish shimmer. 

9. Zinc.—The following gold test is very sensitive: To a few 
drops of a dilute gold solution add a few drops of arsenic acid, two 
or three drops of ferric chloride solution, and two to three drops of 
hydrochloric acid. Dilute the mixture to 100 cc., and drop in a piece 
of zinc. Around the zinc the solution assumes a purple color, which, 
by moving the zinc in the solution, is disseminated through it, making 
it appear pink or purple. 

If the solution contains 0.03 mg. of gold, within fifteen minutes a 
beautiful reddish color will be noticed. 

Besides the above reagents many others, such as formaldehyde 
in the presence of alkali, hydrazine sulfate, etc., are capable of reducing 
gold from its solutions. 

10. Hydrogen Sulfide iia Bc in the cold, black gold disul- 
fide from gold solutions: 


2Aut ¥% +43HS —J 6Ht+ S +AuoSo. 


Gold disulfide is insoluble in acids; but is readily soluble in aqua 
“regia, forming auric chloride, AuCls. 


The disulfide is difficultly soluble in ammonium sulfide, but more | 


readily soluble in potassium polysulfide, forming a sulfo salt: 
Aur2S2+ K2Se = 2K AuSo. 


From this solution hydrochloric acid precipitates a yellowish- 
brown sulfide: 
2KAuSe+2H* ae 2K*+ H2S +Au2S3(?). 





* Vanino and SEEMANN, Ber., 1899, 1968. 
+ Pharm. Chem. Centralbl., 27, 321. 


Soe 


— of a. eT 


GOLD 265 


From a hot solution, hydrogen sulfide precipitates brown, metallic 
gold: 
S8Au’ +++3H2S+12H20 — 30H*+3804"- + 8Au. 


The finely divided metallic gold is soluble in hot potassium or 
sodium polysulfide, forming a thio salt: 


2Au+ KoS4 <—F 2KAuSo. 


On account of its softness, gold is always alloyed with silver and copper 
when used for coins or for jewelry. If such an alloy is treated with nitric 
acid, the copper and silver are dissc!ved and the gold usually remains as a 
brownish powder. ‘To prove the presence of gold, filter through a small filter, 
dry the filter, roll up the paper and wrap a platinum wire around it. Set 
fire to the paper and allow it to burn quietly.. The ash must not be too 
strongly heated, for the gold would then melt and alloy with the platinum 
wire. Melt the ash with a little sodium carbonate on a charcoal stick; a gold 
button forms with the characteristic yellow color. The gold button can be 
pressed into a leaf in the agate mortar, transferred to a watch-glass, and 
dissolved in a little aqua regia, forming auric chloride. Carefully evaporate 
the solution to dryness, dissolve the residue in a little water, add a dilute solu- 
tion of stannous chloride, and the presence of gold is shown by the formation 
of the purple of Cassius. The hydrogen peroxide or zinc tests are still more 
delicate (see p. 264). 

If it is a question of detecting very small amounts of gold (as in the case 
of many copper coins), the above method is unsuitable. In such cases, 
extract the gold and silver by melting with lead, and remove the lead by cupel- 
lation. Melt 5 to 10 gms. of the auriferous copper (or more in some cases) 
with 120 gms. of pure lead in a flat dish of infusible stone (a scorifying-dish), 
in a muffle with access of air. The copper and a part of the lead are oxidized, 
and the oxide unites’ with the silica of the dish to form a readily fusible slag, 
which eventually covers the unaffected lead and the dissolved silver and 
gold. This operation is known as scorification. When this point is reached, 
pour the molten mass into an iron scorification pan, previously well chalked. 
As soon as the mass becomes cool, remove it from the pan, and hammer the 
slag from the enclosed lead button and weigh the latter. Then put it on a 
cupel (a sort of crucible made of bone ash), of about the same weight as the 
lead button or a little heavier. Place the cupel in the muffle and again heat 
with ready access of air. The lead melts and is oxidized; the resulting lead 
oxide melts at 980° C. and is partly absorbed by the porous eupel and partly 
volatilized, leaving a kernel of silver and gold behind. Tlatten the metallic 
kernel to foil and treat it with 6-normal nitric acid, which dissolves the silver, 
leaving the gold, usually in the form of powder. Filter off the gold, dry, and 
melt it, as above described, upon the charcoal stick. If the alloy of gold and 
silver (obtained after cupellation) contains three parts of silver to one part of 
gold, the gold remains after separation with nitric acid as a thin-as-paper, coher- 
ent, brownish mass (which becomes hard on igniting), with the characteristic 
gold color. If the proportion of silver to gold is greater than 3 : 1, the sepa- 
ration by means of nitric acid will be complete and the gold will be left as a 
powder, If the ratio of silver to gold is less than 3 : 1, the separation by means 


266 HYDROGEN SULFIDE GROUP 


of nitric acid is incomplete, and the gold residue usually appears yellow, and 
still contains some silver. In‘this case, add more silver with 1 gm. of lead 
and subject the mixture once more to cupellation, when the subsequent sepa- 
ration by means of nitric acid will be complete. 

In order to detect very small amounts of gold in ores, a similar procedure 
is used. 

If one does not possess a muffle furnace, the tiresome wet process must 
be used. For example, if it is desired to detect the presence of gold in pyrites, 
roast a large amount of the ore in the air until all the sulfur has been burned 
off, then treat with bromine water, and allow to stand twelve hours. Filter 
the solution (which now contains all of the gold as auric bromide) and boil 
off the excess of bromine. Add ferrous sulfate and a little sulfuric acid, boil 
the solution again and filter through a small filter. Wash the residue on the 
filter, dry and then melt it on the charcoal stick. | 

According to these methods a fractional part of a milligram of gold can be 
detected with certainty. 


REACTIONS IN THE DRY WAY 


All compounds of gold, when heated with soda on the charcoal 
stick, yield a malleable, metallic button, soluble only in aqua regia. 
The solution in the latter reagent should be evaporated, the residue 
dissolved in water and tested with stannous chloride, hydrogen per- 
oxide, or zinc. 


PLATINUM, Pt. At. Wt. 195.2 
Sp. Gr. =21.48. M. Pt.=1755° C. 


Occurrence.—Platinum is found free in nature, usually accom- 
panied by the other so-called platinum metals. 

Metallic platinum is grayish white; in a finely-divided state it is 
grayish black. The metal is not ordinarily attacked by mineral acids; * 
it dissolves in aqua regia, forming chloroplatinic acid, He[PtCle] (not 
platinum chloride, PtCl4). If, however, the platinum is alloyed 
' with silver, provided sufficient silver is present, it dissolves in nitric 
acid, forming a yellow solution. Like tin, platinum forms two 
oxides: platinum monoxide, PtO; and platinum dioxide, PtOc. 

Both oxides may be obtained by the careful ignition of the cor- 
responding hydroxides. They are exceedingly unstable, being decom- 
posed by gentle ignition into metal and oxygen; all the remaining 
platinum compounds behave similarly. | 

The most important of the platinum compounds are the chlo- 





* By boiling concentrated sulfuric acid in platinum dishes 10 cc. of acid will 
dissolve 3.8 mg. of platinum in ten minutes. Le Roy and McCay, 8th Internatl. 
Cong. App. Chem. 1, 351 (1912). 


PLATINUM 267 


rides. By dissolving platinum in aqua regia, chloroplatinic acid is 
always obtained, from which the di- and tetrachlorides may be 
derived; these unite with hydrochloric acid to form the complex acids 


PtCl4a+2HCl = Ho2[PtCle] (Chloroplatinic acid—orange red crystals), 


PtCle+2HCl=Hoe[PtCl4] (Chloroplatinous acid—known only in solu- 
tion). | 


The aqueous solution of chloroplatinic acid is yellowish orange; 
a solution of chloroplatinous acid, containing the same quantity of 
platinum, is dark brown. 

The potassium and ammonium salts of chloroplatinous acid are 
soluble in water; the corresponding salts of chloroplatinic acid are 
difficultly soluble in water and insoluble in 75 per cent alcohol. 


REACTIONS IN THE WET WAY 


A solution of chloroplatinic acid should be used for these reactions. 
1. Ammonium and Potassium Chlorides produce in concentrated 
solutions yellow precipitates (cf. pp. 78 and 89): 


[PtClg)"+2K+t > Ko[PtCle]; [PtCle]- +2NH4t > (NH4)2[PtClo]; 


Both salts are difficultly soluble in water, but practically insoluble 
in 75 per cent alcohol and in concentrated solutions of potassium 
and ammonium chlorides. This last property is utilized in separating 
platinum from gold and other metals. 

2. Alkali Iodides give a brownish-red coloration due to the forma- 
tion of the less ionized [PtI¢]~- 


[PtCle]-+6I- — [PtIe¢]—-+6ClI-. 


3. Hydrogen Sulfide precipitates in the cold very slowly, but 
quickly on warming, dark-brown platinum disulfide: 


H.2[PtCle]+2H2S =6HCI1+PtSo. 


Platinum sulfide is insoluble in mineral acids, but readily soluble 
in aqua regia. It is difficultly soluble in alkali sulfides, but more 
readily soluble in alkali polysulfides, forming a thio salt, which is 
decomposable by acids, with precipitation of platinum sulfide. 

4. Ferrous Salts do not reduce chloroplatinic acid in the pres- 
ence of acids (difference from gold), but cause precipitation of all 
the platinum (on warming) in a solution which has been neutralized 
with sodium carbonate: 


268 HYDROGEN SULFIDE GROUP 


[PtCle]- +6CO37+4Fet ++6H20 — 6CO2 T +6CI-+4Fe(OH)3+Pt. 


5. Oxalic Acid does not precipitate platinum (difference from 
gold). 


6. Formic Acid precipitates from neutral, boiling solutions all the . 


platinum in the form of a black powder: 
H2[PtCle] +2HCOeH =6HCI+2CO02+Pt. 


Formic acid 


An acid solution must be neutralized with sodium carbonate before 
making this test. 

7. Stannous Chloride reduces chloroplatinic acid to chloroplatin- 
ous acid only, not to metal: 


[PtCle]-+Snt + =Snt*+t+++2Cl-+[PtCh]-. 


8. Glycerol and Sodium Hydroxide reduce chloroplatinic acid on 
warming, with the separation of black, pulverulent metal: 


C3Hs(OH)3+8[PtCle]> +16OH- — 18Cl +CO37 +C2047 +12H20 +8Pt. 


9. Carbon Monoxide, on being passed into a solution of chloro- 
platinic acid containing sodium acetate, colors the solution a beau- 
tiful red owing to the formation of colloidal platinum (difference from 


palladium). After standing some time all the platinum is precipitated a , 


as a black powder and the supernatant liquid is colorless: 


H2[PtCle] +2CO+6C2H302°+2H20 -6Cl-+6HC2H302+ 2COe2T +Pt. 


10. Zinc, cadmium, magnesium or aluminium reduces chloroplatinic 


acid to metal: 
H2[PtCls]+3Zn — 3ZnCle+He ft +Pt. 


The precipitated metal is in such a finely divided condition that it 
tends to run through the filter, especially on being washed with pure 
water; by washing with salt solution a clear filtrate can be obtained. 


11. Hydrazine Hydrochloride, NoH4-2HCl, readily reduces chloro- ; ’ 


platinic acid in ammoniacal solutions; the platinum is deposited 
partly as mirror upon the sides of the vessel containing the solution. 


(NHa4)2[PtCle]+N2H4-2HCI+6NH3 — 8NH4Cl+Ne J +Pt. 


12. Formaldehyde in alkaline solutions precipitates the platinu 
as extremely- finely-divided platinwm black: 


He[PtCl6]| +HCHO+60H- — COz fT +6CI”-+5H20+Pt. 


The precipitate may be washed with alkali salt solutions, but with pure © 


water a black colloidal solution of platinum is formed. 


ee ee a ee ee ne en ee eee 





: 


PLATINUM 269 


Preparation of Chloroplatinic Acid for Use as a Reagent 


Since chloroplatinic acid is used not only for the qualitative detection of 


- ammonium and potassium, but also for their quantitative separation, methods 


for preparing a solution of this reagent will be described. 

1. Preparation of Chloroplatinic Acid from Metallic Platinum.—Almost 
all commercial platinum contains iridium; and although pure iridium is prac- 
tically insoluble in aqua regia, it dissolves considerably in this reagent if it 


is alloyed with platinum. Moreover, platinum forms with aqua regia not 


only cholorplatinic acid, but also chloroplatinous acid (the most harmful 
of all impurities for this reagent) and nitroso-platinic chloride, [PtCl](NO)s. 
These facts must be borne in mind in preparing the reagent. 

First of all, clean the strips of platinum by boiling them with concentrated 


hydrochloric acid and washing with water. Then place the platinum in a 


capacious flask, pour over it concentrated hydrochloric acid, and add nitric 
acid little by little, with continuous, gentle heating on the water-bath. All 
the platinum and some iridium are thus brought into solution, while some 
of the latter metal usually remains undissolved as a black powder. 

Decant the solution (without stopping to filter) into a porcelain evapo- 
rating-dish and evaporate to sirupy consistency. Dissolve the residue in 
water, add some sodium formate and sodium carbonate until the solution is 
slightly alkaline. Heat the solution to boiling, which causes the precipitation 
of the platinum and iridium in a few minutes as a black powder. This opera- 
tion should be performed in a large evaporating-dish, on account of the violent 
effervescence due to the escape of carbon dioxide. Pour off the supernatant 
liquid and wash the residue several times with hydrochloric acid to remove 
the sodium salt, and finally with water to remove the acid. Dry the powder, 
which contains iridium and platinum (in the presence of one another, but 
not alloyed together) and ignite it strongly in a porcelain crucible over the 
blast-lamp (whereby the iridium is made insoluble in aqua regia), and 
then weigh. Dissolve the ignited gray metal (at as low a temperature as 
possible) in hydrochloric acid, with gradual addition of nitric acid. Consider- 
able quantities of nitroso-platinic chloride are formed by this operation. On 


evaporating with water, this compound is decomposed into chloroplatinie 


acid, with evolution of oxides of nitrogen: 
[PtCl.](NO).+H,0=NO, T +NO f +H,[PtCl,J. 


As, however, a part of the NO. (or N.O,) remains in solution, some more 
nitric and nitrous acids are formed by the action of water, 


N20, + H,O = HNO; -}- HNO,, 


which yields nitrosyl chloride with the hydrochloric acid present, and causes 
the formation of more nitroso-platinic chloride. 

It is necessary, therefore, to evaporate the solution alternately with hydro- 
chloric acid and with water until no more nitrous fumes are given off. The 
solution thus obtained always contains some chloroplatinous acid, and is 
intensely brown. In order to change this last compound into chloroplatinic 
acid, saturate the warm solution with chlorine gas (whereby its color becomes 
much lighter) and then evaporate (at as low temperature as possible) till it 


270 HYDROGEN SULFIDE GROUP 


becomes of sirupy consistency. After cooling, the sirup crystallizes to a 
yellowish-brown mass, which may be dissolved in a little cold water, and the 
insoluble iridium filtered off. 

If there is a considerable amount of the latter metal, ignite it in a pongelaan 
crucible and weigh. The weight of the iridium should be deducted from the 
prev ious weight of the mixture, in order to find out how much platinum re- 
mains in solution. 


Dilute the filtered solution with water until 100 ec. of the solution contains 


10 gms. of platinum. 
2. Preparation of Chloroplatinic Acid from Platinum Residues.—These 
residues consist of potassium chloroplatinate and the alcoholic wash-waters. 
By evaporating an alcoholic solution of chloroplatinic acid, chloroplatinous 
acid and ethylene are formed, which yield ethylene platinous chloride: this last 
compound gives no precipitation with potassium or ammonium salts: 


H.[PtCl]+2C,.H;OH =CH;CHO+4HCI+C,H,[PtCl.]4+ H.0. 
Alcohol Aldehyde 


On evaporating an alcoholic solution of this soluble organic platinum 
compound, it is changed into an insoluble powder which is explosive when dry, 
insoluble in acids, and completely decomposed by strong ignition only. 

To separate platinum, therefore, from platinum residues, first evaporate 
the alcoholic solution to dryness, take up the residue in water, and pour the 


solution into caustic soda (sp.gr. 1.2), to which 8 per cent of glycerol * has 


been added. Heat the liquid to boiling, which causes the platinum to be pre- 
cipitated as a heavy black powder, 


FOO. GHP OCU HO = 36HCI+ 2CO.+2H.2C.0.+6Pt. 


Glycerol Oxalic acid 


Wash the powder first with water, then with hydrochloric acid, and finally 
with water again. Dry, ignite (to destroy any of the organic compound), 
weigh and transform, as before, into chloroplatinic aicd. 


REACTIONS IN THE DRY WAY 


All platinum compounds, when heated with soda on charcoal, 
are reduced to the gray, spongy metal, which assumes a metallic luster 
on being rubbed with a pestle in an agate mortar. It can be dis- 
tinguished from gold by its color, and from lead, tin, and silver by its 
infusilibity and insolubility. 


Separation of Gold from Platinum 


Precipitate the platinum with a solution of ammonium chloride; 
filter and treat the solution with ferrous sulfate to precipitate the 
gold. 





*Z. anal. Chem., 28, 509. 


a ee ca el i i 


ANALYSIS OF GROUP II 271 


Analysis of Group II 


Gold and platinum are two well-known metals which really belong to this 
group, but they are seldom present in large quantities in ordinary mineral 
analysis and the detection of small quantities is so important that it is custom- 
ary to test for them separately as has been indicated (cf. pp. 265 and 270). 

The metals of the second group are precipitated as sulfides by means of 
hydrogen sulfide in the presence of 0.3-normal mineral acid. The theory 
governing the precipitation of sulfides was discussed on pp. 13, 47, 145, 186, 
etc. In 0.3-normal mineral acid, the concentration of sulfur ions, formed by 
the ionization of hydrogen sulfide, is made so small that the solubility products 
of the sulfides of zine, nickel, cobalt, iron and manganese is not reached unless 
large quantities of these elements are present in solution. In quantities up 
to 0.5 gm. of metal, zine sulfide is the only one which is at all likely to pre- 
cipitate in acid of this concentration, and then only when it is carried down 
with a considerable quantity of some second-group metal. The solubility- 
product of lead sulfide is about 4.2<107-28 and that of cadmium sulfide is 
3.6X10~*9, while that of zine sulfide is 1.210-?%. Compared with mercuric 
sulfide, with its solubility-product of 4.010753, lead sulfide is quite soluble, 
but compared with zine sulfide, it is much less soluble. In precipitating these 
sulfides there is some tendency for the sulfide precipitate to carry down 
with it some of the metals of the succeeding groups partly by adsorption, but 
this tendency is so slight that even with zinc the test is usually obtained in 
the proper place when as much as 2 mgms. is originally present. 

The solubility of the sulfides of the second group varies so greatly that it 
is possible to separate them from one another by regulating the acid con- 
centration so that enough sulfur ions are furnished by hydrogen sulfide to 
precipitate one metal and not another. For the same reason, when hydrogen 
sulfide is passed into the acid solution, the most insoluble sulfide is precipitated 
first and very little, if any, of a more soluble sulfide is formed until the precip- 
itation of the former is complete. In this respect the behavior of arsenic forms 
an apparent exception; but this is due, as already pointed out (p. 230), to the 
absence of an appreciable quantity of arsenic cations in the solution of an 
arsenate. The order in which the metals are precipitated as sulfides from 
cold solutions, as the acid concentration is progressively decreased, is approx- 
imately as follows: arsenic, mercury and copper; antimony, bismuth and stan- 
nic tin; cadmium, lead and stannous tin; zinc, iron, nickel, cobalt and man- 
ganese. It will be noticed that zinc is placed a little in front of nickel and 
cobalt in this arrangement. This is the proper order for the precipitation, but 
after the sulfides have been formed a short time, nickel and cobalt sulfides 
become less soluble than freshly-precipitated cadmium, lead or stannous 
sulfide. 

The theory of the solution of the sulfides has also been indicated (pp. 145, 
181). The more soluble sulfides can be dissolved by merely increasing the con- 
centration of the hydrogen ion. Thus the sulfides of antimony, tin, lead, cad- 
mium, etc., can be dissolved by treatment with concentrated hydrochloric 
acid.. Heating the solution to expel hydrogen sulfide hastens the dissolving, 
but the most effective expedient is to add an oxidizing agent, which oxidizes the 
sulfur ions as fast as they are formed in solution. Thus even mercuric and 
arsenic sulfides will dissolve readily on being treated with aqua regia or with 
bromine water, 


272 HYDROGEN SULFIDE GROUP 


The second group of metals may be divided into two smaller sub-groups. 
The first of these contains mercury, lead, bismuth, copper and cadmium and 
represents those elements whose sulfides are not dissolved readily by ammonium 
polysulfide. Most of these sulfides show a slight tendency to dissolve in the 
polysulfide and the tendency is most marked with the sulfides of mercury 


and copper. From 5 to 10 mgms. of CuS and 0.5 to 1.0 mgms. of. HgS may dis- 
solve in the polysulfide when the original substance contains considerable of — 
these elements. None of these sulfides dissolve to any extent in ammonium ~ 


monosulfide. 

The other subgroup contains arsenic, antimony and tin. The sulfides 
of these elements dissolve readily in 6-normal ammonium polysulfide solution. 
It is easy to dissolve 0.5 gm. of arsenic either as As»S; or as As»S;, 0.5 gm. of 
antimony as Sb.S; or 0.5 gm. of tin as SnS, in ether ammonium monosulfide 
or the polysulfide. Scarcely any SnS and only about 0.1 gm. of antimony as 
Sb.S; will dissolve in 25 ce. of 6-normal ammonium monosulfide, but as much 
as.0.5 gm. of either SnS or Sb.S; will dissolve in the same quantity of am- 
monium polysulfide. Ammonium monosulfide dissolves scarcely any of the 
sulfides of the other subgroup, so that when the antimony and tin are known 
to be present in the higher state of oxidation, it is advisable to use ammonium 
sulfide rather than the polysulfide for the separation of the two subgroups. 


TasBLeE VII.—SEPARATION OF THE COPPER AND TIN GROUPS 





Solution may contain: Hgt+, Pbtt, Bitt++, Cutt, Cdtt, Astt+t, AsO;=, AsOg, 
Sbt+++, Sbht++++, SbO=, Sntt, Snt++** and Groups III, IV and V. Make the 
solution 0.3-normal with HCl and saturate with H»S. Filter and examine the filtrate 
for Groups IIIT, IV and V. Treat the precipitate with (NH4).Sz. (1) 





Residue: HgS, PbS, Bi.S;, | Solution: [AsS,]=, [SbS,]=, [Sn83]>. 
CuS, CdS. Add HCl. (3) 
Examine by Table VIII. (2) 





Precipitate: As.S;, SbS;, | Filtrate: NH,Cl. 
Sn82: Reject. 
Examine by Table IX. (3) 








PROCEDURE 


1. Add to the neutral solution 5 ec. of 6-normal hydrochloric acid and 
saturate with hydrogen sulfide in the cold. Filter promptly and wash with 
hot water containing hydrogen sulfide. Heat the filtrate to boiling and pass 
hydrogen sulfide into it for ten minutes to make sure that arsenic will 


be precipitated. If a yellow precipitate is formed slowly by this treatment, ~ 


filter it off through a new filter and evaporate the filtrate nearly to dryness. 
Then add 10 ce. of 12-normal hydrochloric acid, saturate the cold solution 
with hydrogen sulfide, heat nearly to boiling and again introduce hydrogen 
sulfide. Cool, dilute to 100 cc. and saturate again with hydrogen sulfide. 
Filter off the precipitate and examine the filtrate for the metals of Groups 
III, IV and V (cf. p. 189 or 192).. This repeated treatment with hydrogen 
sulfide is necessary when considerable arsenic is present (cf. p, 230). 


Pe eee 


7 


ANALYSIS OF GROUP II 273 


2. Transfer the precipitated sulfides to a porcelain dish and add about 
10 cc. of ammonium polysulfide (or, better, ammonium monosulfide if it is 
certain that the precipitate contains no stannous sulfide nor a large quantity 
of antimony trisulfide). Cover the dish and warm gently for a short time with 
frequent stirring. Dilute with a little water, filter and wash with hot water 


containing a little ammonium sulfide. If the original residue was large and 


the treatment with ammonium sulfide has evidently reduced its bulk, it is advis- 


+ able to repeat the treatment with ammonium sulfide in order to make sure that 
all of the arsenic, antimony and tin sulfides have been converted into soluble 


thio salts. Examine the residue by the method outlined in Table VIII. 

3. Dilute the ammonium sulfide extract with an equal volume of water, 
add hydrochloric acid until the stirred solution is acid to litmus, heat to 
boiling and filter as soon as the sulfides of arsenic, antimony and tin have 
settled. Reject the filtrate and examine the precipitate by the method out- 
lined in Table IX. 


TaBLE VIII.—ANALysIs OF THE CopPpER GROUP. 





Residue from Table VII: HgS, PbS, BiSs, CuS, CdS. Boil with 2-normal HNO3. (1) 





Residue: HgS. | Solution: Pb++, Bit++, Cutt, Cdt++. Add 6-normal H2SOs, 
Dissolve in evaporate, dilute and filter. (8). 
aqua regia or 
in bromine | precipitate: | Filtrate: Bit++, Cut+, Cd++. 
water. Di-| ppso,. Add 6-normal NH,OH in excess. — (5) 
oe a bil Dissolve in 
ye | NAGC2H02 Precipitate: Solution: 


ae ed ee. Bi(OH);. Dissolve | [Cu(NHs)4]*++,[Cd(NHs),]++. 
itateof HgoCl. 7 Il 24 | in HCl, concentrate | A blue solution shows Cu. 
or agray pre- en0wW Pre | the solution and | If in doubt, add HC.H;O, and 
cipitate of eno | pour it into consider-| K,Fe(CN)g to a portion of the 
HgCh + Hg| «p mt able water. White| solution. Red precipitate 
shows _pres- eye re. precipitate of BiOCl| shows Cu. (7) . 
ence of Hg. i of Pb.| indicates the pres- | Decolorize the solution if neces- 
(2) (4) ence of bismuth. sary with KCN and test with 
Filter and add| HS. Yellow precipitate - 
NaxSnO,. Black| shows presence of Cd. (8) 
residue is Bi. (6) 




















PROCEDURE 


1. Treat the residue from the ammonium sulfide treatment (Table VII) 
with 2-normal nitric acid, boiling for a short time. Filter and wash with hot 
dilute nitric acid. Examine the filtrate by (38). 

2. Transfer the residue, and a part of the filter if necessary, to a porcelain 
dish and digest with hot aqua regia, or warm with saturated bromine water. 
Boil off the excess of chlorine or bromine, but do not evaporate to dryness on 
account of the danger of losing some mercuric salt by volatilization. Dilute 
somewhat and filter. Sometimes, when the elements of the copper group 
are present in large quantity, the residue of sulfur (and filter paper) will con- 


274 HYDROGEN SULFIDE GROUP 


tain metastannic acid. To recover the tin, digest the residue with bromine 
again if it is at all dark colored, filter and reject the filtrate. Then warm 
the_residue with a little ammonium sulfide solution and add it to the solu- 
tion obtained upon treatment of the original hydrogen sulfide precipitate with 
ammonium polysulfide (Table VII). Test the solution obtained by the above 
treatment with aqua regia, or bromine, for mercury by adding stannous chloride 
solution. <A white precipitate of H. gnCle or Hg.Br2 shows that mercury is present. 
The precipitate turns gray when an excess of stannous chloride is added 
(cf. p. 200). 

3. To the filtrate obtained in (1) add 10 cc. of 6-normal sulfuric acid and 
evaporate in a porcelain dish until white fumes of sulfuric acid are evolved. 
The purpose of this treatment is to expel all nitric acid, in which lead sulfate is 
somewhat soluble. Cooland pour into 10 ce. of water, rinsing out the dish with 
a little cold water. Stir well and let the solution stand five minutes but not much 
longer. Filter and treat the fitrate by (5). When much bismuth is present 
some basic bismuth sulfate is often precipitated with the lead sulfate, but the 
greater part of the bismuth will remain in solution, particularly if the water 
is kept cold. The basic bismuth sulfate, (BiO)»SO,, is more coarsely crystalline 
than lead sulfate. When such a precipitate is obtained, treat it with 10 ce. 
of hot 6-normal hydrochloric acid, add 10 cc. of 6-normal sulfuric acid and evapo- 
rate as-above. The precipitate of lead sulfate now obtained will be free from 
bismuth. The treatment of the first sulfate precipitate with hydrochloric 
acid will dissolve a part or all of the lead sulfate. 


4. Dissolve the precipitate of lead sulfate by pouring small portions of 


hot 3-normal ammonium acetate solution through the filter. Do not use over 
25 cc. of the acetate solution and wash once with hot water after each addition 
of the acetate. To thesolution thus obtained add a few drops of potassium 
chromate solution and 2 to 5 cc. of 6-normal acetic acid. A yellow precipitate 
of lead chromate, insoluble in acetic acid, shows the presence of lead. Bismuth 
chromate is readily soluble in acetic acid. 

5. Neutralize the filtrate obtained in (3) with ammonia, using a liberal 
excess. This serves to precipitate bismuth hydroxide and form soluble com- 
plex cations with copper and cadmium. If the original hydrogen sulfide pre- 
cipitate was not washed thoroughly, a precipitate of ferric hydroxide or of 


aluminium hydroxide may be obtained at this point. Both of these pre-. 


cipitates have a different appearance tiian bismuth hydroxide and do not give 
the confirmatory test for bismuth. Examine the filtrate by (7). 

6. Pour a little 6-normal hydrochloric acid through the filter containing 
the bismuth hydroxide and evaporate till only a few drops of liquid remain, 
or a moist residue. Add 1 or 2 cc. of water and pour into 100 cc. of hot water. 
After standing two or three minutes, filter off any BiOCl and wash once with 
cold water. Prepare a fresh solution of sodium stannite by taking a few drops 
of stannous chloride solution, diluting with 5 cc. of water and adding sodium 
hydroxide solution, a few drops at a time until the Sn(OH), which first forms 
redissolves. Pour this solution through the filter containing the BiOCl A 
black residue of bismuth is obtained when this element is present (cf. p. 213). 

7. If the ammoniacal filtrate from (5) is blue, copper is present. The only 
other element which is likely to be confused with copper at this point is 
nickel, which also forms a blue solution with ammonia. The color of the 
ammoniacal nickel solution is very much paler than that of the corresponding 


™ 


ANALYSIS OF GROUP II 275 


copper solution, and it is rare that enough nickel is left with the hydrogen 
sulfide precipitate, due to incomplete washing, to cause trouble. In cases 
of doubt, acidify a little of the ammoniacal solution with acetic acid and add 
_ one drop of potassium ferrocyanide solution. <A red precipitate of Cue[Fe(CN)c], 
is formed if copper ts present. The test is much more sensitive than the blue 
test with ammonia; nickel, under the same conditions, gives a less character- 
istic green precipitate. 

(8) If the ammoniacal solution is blue, add potassium cyanide solution 
until it is colorless, or only a few drops if the solution is already colorless. Pass 
hydrogen sulfide into the solution for half a minute. A yellow precipitate of 
cadmium sulfide is formed if cadmium is present. A red precipitate of 
(CSNH:)2 may be formed if too much H,S is added (p. 219). 

Sometimes a small black precipitate is obtained in the cadmium test which 
is most likely due to a little mercury or lead that was not properly removed 
from the solution. In such cases, filter off the precipitate, wash it with water 
until free from cyanide and boil gently for five to ten minutes in a covered 
dish with about 15 cc. of 1.2-normal sulfuric acid (6-normal acid diluted with 
_ four volumes of water). This serves to dissolve cadmium sulfide, but will not 
dissolve copper or mercury ‘sulfides and should convert lead sulfide into lead 
sulfate. Filter, dilute with three times as much water and saturate with 
hydrogen sulfide. A yellow precipitate of cadmium sulfide will now be 
obtained if cadmium is present. 


TABLE IX.—ANALYSIS OF THE ARSENIC-TIN GROUP 





Precipitate from Table VII: As.S;, SboS;, SnS2. Warm with 12-normal HCl. (1) 





Residue: As.S;. Dissolve in 6- 
normal HCl and _ KClIOs. 


Solution: Sbt++, Snt++++, Evaporate to small 
volume, pour upon clean platinum foil, and place 


Evaporate, dilute, neutralize 
with NH,OH and add MgCl. - 
NH.Cl solution. A white pre- 
cipitate of MgNH,AsO, indi- 
cates As. Dissolve in 6-nor- 
mal HCl, and treat with H2S. 
Yellow precipitate of AsS5 


a clean piece of zinc in the solution; a black spot 
on the platinum indicates Sb. When the evolu- 
tion of hydrogen has ceased, remove the zine and, 
if any tin deposit adheres to the platinum, rub it 
off and dissolve it, with the remaining zinc, in a 
small test-tube in one or two drops of concentrated 
hydrochloric acid. Dilute with water and add a 


or AsS3 shows As. (2) few drops of mercuric chloride solution. A white 


or gray precipitate shows Sn is present. (3) 








PROCEDURE 


1. Transfer the ‘sulfides of arsenic, antimony and tin, which are mixed 
with considerable free sulfur, to a porcelain dish and digest for about fifteen 
minutes with 12-normal hydrochloric acid. Heat gently, but do not boil 
hard. This serves to dissolve antimony pentasulfide (forming antimony tri- 
chloride) and stannic sulfide (forming stannic chloride) but does not dissolve 
much if any arsenic pentasulfide. Dilute with a little water and filter. Treat 
the solution by (8). 

2. Heat the residue of arsenic sulfide with 15 ec. of 6-normal hydrochloric 
acid and add powdered potassium chlorate from time to time in small portions. 


276 HYDROGEN SULFIDE GROUP 


When all the yellow sulfide has dissolved, dilute somewhat and filter off the ‘3 


‘ 


residual sulfur. Evaporate the solution to about 2 cc., add 6-normal ammonia 
in liberal excess and a few drops of magnesium-ammonium chloride reagent. 


If no precipitate forms at once, stir vigorously and rub the sides of the glass 
vessel with the stirring rod and let the solution stand at least an hour. Ifa _ 
precipitate of magnesium ammonium arsenate, MgNH,AsQ,, forms, pour — 
off the liquid through a filter and dissolve the precipitate in 10 ce. of 6-normal — 


hydrochloric acid. Saturate the solution with hydrogen sulfide, heat to 

boiling and again introduce hydrogen sulfide for at least five minutes. A yellow 

precipitate of arsenic sulfide will form if as much as 1 mg. of arsenic is present. 
3. Evaporate the solution of SbCl; and SnCl,, obtained in (1), to a very 


small volume and pour a few drops of the concentrated solution upon a piece © 


of clean platinum foil. Place a small piece of bright zinc in the liquid so that 
it comes in contact with both solution and platinum. After a few seconds, 


take away the zinc and notice whether a coal black spot of antimony has formed — 


upon the platinum. Mercury and copper if present will also be precipitated 
_ upon the platinum, but the mercury deposit is gray and the copper deposit 





a 





is red. ‘The antimony spot is insoluble in sodium hypochlorite solution 


(cf. pp. 235, 249). 

Replace the zinc upon the platinum foil and allow the action to continue 
until there is no more evolution of hydrogen. Then rinse carefully with dis- 
tilled water, taking care not to disturb the contact of the zinc with the platinum. 
Take away the zinc and carefully remove any tin that may adhere to the plati- 
num. Dissolve the mixture of tin and zinc in one or two drops of concentrated 


hydrochloric acid. Dilute the solution in a very small test-tube with a little — 


water and add a few drops of mercuric chloride solution. A white precipitate 
of mercurous chloride which turns gray shows the presence of tin. 


SUPPLEMENTARY PROCEDURES 
Analysis of Arsenic—Tin Group 


A. If the sulfide precipitate, obtained by acidifying the ammonium 
polysulfide solution, consists chiefly of arsenic pentasulfide, it is best to 
dissolve the arsenic sulfide first by means of ammonium carbonate solution 
(p. 230). Then the residue of Sb.S;, SnS. and § can be examined for antimony 
and tin. 


To detect the arsenic in the ammonium carbonate solution, add hydro- 


chloric acid, which reprecipitates arsenic pentasulfide. Dissolve the sulfide 
as indicated in Table X and confirm the arsenic test as described there. 

B. An excellent method for separating antimony and tin is based upon the 
fact that antimony trisulfide is less soluble than the sulfides of tin. To accom- 
plish the separation, proceed as follows: 


Dissolve the sulfides of antimony and tin as described above in exactly — 


10 ec. of 12-normal hydrochloric acid; when there is no further evolution of 
hydrogen sulfide gas, dilute with 3 cc. of water and filter with the aid of suc- 
tion. Dilute the filtrate to exactly 55 cc. and transfer the solution to a small 
flask. Place the flask in a beaker of boiling water and introduce hydrogen 
sulfide into the solution in the flask for ten minutes but no longer. Filter 
and wash the precipitate with hot water. 

An orange precipitate of antimony trisulfide will be formed when only 


a eS — Fy eh 


i es 4 A * * 
ae Se a ae 


ANALYSIS OF GROUP II | 277 


1 mg. of Sb is present. Dissolve the precipitate in a little 12-normal hydro- 
_ chlorie acid, evaporate nearly to dryness, dilute and test with platinum and 
zine as described above. 

The filtrate from the Sb.S; can be used for the tin test. Instead of reduc- 
ing the tin with zinc, metallic lead can be used (p. 254). Evaporate the solu- 
tion to about 3 cc., but not to dryness on account of the danger of losing 
_ stannic chloride by volatilization, add 10 cc. of water, 10 gms. of granulated 
lead and heat gently for ten minutes. Pour the hot solution through the 
filter into 10 cc. of 0.2-normal mercuric chloride. Filter off any mercurous 
chloride that may be precipitated, dissolve any adhering lead chloride by wash- 
ing with hot water and, if not already blackened, add ammonia to the precipi- 
tate (cf. p. 201). 


GROUP I. HYDROCHLORIC ACID GROUP 


To this group belong silver, mercury (in mercurous compounds), 
lead, thallium, and (under some circumstances) tungsten. 


SILVER, Ag. At. Wt. 107.88 
Sp. Gr.=10.5. M. Pt.=961° C. 


Occurrence.—Silver occurs both native and combined (chiefly with 
sulfur, arsenic, and antimony). 

Of the silver-bearing minerals the following may be mentioned: 
horn silver, AgCl; argentite, AgeS; pyrargyrite, AgsSbS3; and prous- 
tite, AgsAsSz3. 

Silver is also found with tetrahedrite, and with galena. 

Metallic silver is of a pure white color. It absorbs oxygen in the 
molten state, which it gives up (with tiny explosions) on cooling. 

The proper solvent for silver is nitric acid. It is insoluble in 
dilute hydrochloric and sulfuric acids, but dissolves readily in boiling 
sulfuric acid, with evolution of sulfur dioxide: 


Ago+2H2SO4 = 2H20+S802+AgeSOz. 


The solubility of silver in concentrated sulfuric acid is utilized 
in separating silver from gold and platinum in alloys. 

Silver forms three oxides: silver suboxide, Ag4sO; silver oxide, 
AgeO; silver peroxide, AgeOo. ' 

Of these oxides, AgeO alone is a basic anhydride; only one series 
of salts is known. 


Silver peroxide is formed at the anode during the electrolysis of a con- 
centrated, aqueous solution of silver nitrate. The grayish-black crystalline 
powder has the composition: Ag,O,=Ag,0-Ag,0;.* 


Silver oxide, AgeO, is a brownish-black powder, which on being 
heated to 300° is completely decomposed into metal and oxygen. 

Most of the silver salts are colorless; the following, however, are col- 
ored: the bromide (pale yellow), the iodide (yellow), the sulfide (black), 





* BRAUMER and Kuzma, Ber., 1907, 3371. It may be regarded as the silver salt 
of the hypothetical argentic acid, HAgOs; ; its symbol could then be written Biri 
analogous to the aurates, e.g., K[AuO,]. 

278 


SILVER | 279 


the phosphate (yellow), the arsenite (yellow), the arsenate (brown), 
the ferricyanide (orange), and the chromate (reddish-brown). Most 
of the salts are insoluble in water, and are blackened on exposure to 
the light. The nitrete, chlorate, perchlorate, fluoride, nitrite, sulfate, 
and acetate are soluble in hot water. 


REACTIONS IN THE WET WAY 


1. Potassium and Sodium Hydroxides precipitate brown silver 
pxide, 
2Age*+20H- — H20+Age0, 


insoluble in an excess of the precipitant, but readily soluble in nitric 
acid and in ammonia. If the solution in ammonia is allowed to stand, 
black detonating silver is deposited, [AgNH3]20. 

2. Ammonia.—If a neutral solution of a silver salt is cautiously 
treated with ammonia, the first drops produce a white precipitate, 
which changes quickly to the brown oxide, Ag2O. The greater part 
of the silver, however, remains in solution as complex silver-ammonia 
salt; even the oxide is dissolved by an excess of ammonia: 


Ag20+4NH3+H20 — 2/Ag(NH3)2]*+20H-. 


3. Sodium Carbonate precipitates white silver carbonate, which 
becomes yellow on being boiled, being slightly decomposed into oxide, 
with loss of carbon dioxide: 


2Ag*+CO37 => AgeCOs; 
AgeCOz ae Age2O+ COeg T ° 


4. Ammonium Carbonate produces the same precipitate, but 
it is soluble in an excess of this reagent. 

5. Sodium Phosphate throws down in neutral silver solutions 
a yellow precipitate of silver phosphate: 


3Agt+ 2HPOs — HePO4-+AgsPOz. 


Silver phosphate is easily soluble in nitric acid and in ammonia. 
The solution of the phosphate in ammonia is due to the formation 
of complex silver-ammonia ions. 


AgsP04+6NH3 — 3[Ag(NHs3)2]*+PO4=. 


By neutralizing the ammoniacal solution with nitric acid, or the 
nitric acid solution with ammonia, the silver phosphate is repre- 
cipitated. 


280 HYDROCHLORIC ACID GROUP 


6. Hydrochloric Acid and Soluble Chlorides precipitate from s 
neutral and acid solutions white, curdy silver chloride: ‘aa 


Agt+Cl-~— AgCl. 


Silver chloride is appreciably soluble in pure water (ef. p. 21), particu- 
larly on boiling, but it is quite insoluble if an excess of silver nitrate or of 
hydrochloric acid is present (cf. p. 20). 

Silver chloride dissolves to a considerable extent in a large excess of hydro- 
chloric acid or of alkali chloride, due to the formation of complex anions, 
[AgCl;]-, but it is much less soluble in dilute nitric acid. 

It is very soluble in ammonia, 


AgCl+2NH; — [Ag(NH;).]*CI-, 
but is reprecipitated on addition of nitric acid to this solution: 
[Ag(NH;).]+2Ht — 2NH,*+AgCl. 
Silver chloride is also readily soluble in potassium cyanide solution, 
AgC1+2CN~ — [Ag(CN).]-+CI, 
and if such a solution is acidified, silver cyanide will be precipitated; 
Ag(CN),-+H*t — HCN+AgCN; 


In the absence of acid, silver chloride is also readily soluble in a solution 
of sodium thiosulfate: 


2AgCl+ 38,0;> => [Age (S20 3) 3} si of 7 om 


On boiling this solution, silver sulfide is precipitated slowly if considerable _ 
Na.$.O; is present, but more quickly upon diluting. | 

It is evident that the solubility in all the above cases is due to the formation : 
of complex ions containing silver (cf. p. 24). 

Silver chloride is slowly attacked by boiling, concentrated sulfuric acid, 
with evolution of hydrochlori¢ acid, and the formation of crystalline silver 
sulfate, insoluble in sulfuric acid. 

By boiling with caustic soda or caustic potash solution, silver chloride is 
only partially decomposed; in the cold it is unaffected. Sodium carbonate 
solution does not affect it; but by fusing with sodium carbonate it is com- 
pletely decomposed: 


4AgC]1+2Na,CO; =4NaCl+2CO, T +0, fT +4Ag. 


By fusing silver chloride itself a yellow liquid is obtained, which on siti 
solidifies to a horny mass. 


7. Potassium Iodide precipitates yellow, curdy silver iodide, 
practically insoluble in ammonia, but easily soluble in potassium 
cyanide and sodium thiosulfate solutions. | 

The [Ag(NH,).]+ cation furnishes more Ag* cations than silver iodide 


in contact with water, but the [Ag(CN).]~ and [Ag,(8.0;),]>~ are much more 
stable complexes and furnish less Agt cations than silver iodide does. 





SILVER 281 


8. Ferrous Sulfate precipitates gray metallic silver from boiling 
solutions: 


Agt+Fet*— Fettt+Ag. 


Frequently a basic ferric salt is precipitated at the same time, 
particularly from very dilute solutions. : 

9. Zinc, having a greater solution pressure than silver, precipi- 
tates the latter from neutral solutions. Similarly, if insoluble silver 
chloride is covered with dilute sulfuric acid and a piece of zine placed 
in contact with the chloride, the latter will be reduced to metal. 


2AgC1+Zn=Zn*++2Cl-+2Ag. 


10. Hydrogen Sulfide precipitates from neutral, ammoniacal, and 
acid solutions black silver sulfide, 


2Ag+H2S > 2H*+Ag.S, 


insoluble in ammonia, alkali sulfides, and dilute potassium cyanide 
solution. Silver sulfide is so insoluble (cf. p. 21) that enough Agt 
ions are present in aqueous solutions containing [Ag(NH3)e]* or even 
[Ag(CN)2]— to exceed the solubility product of AgeS when in contact 
with slightly ionized hydrogen sulfide. Silver sulfide is, however, 
perceptibly soluble in a concentrated solution of potassium cyanide, 
and easily soluble in hot dilute nitric acid (cf. p. 181). 

11. Potassium Chromate precipitates brownish-red silver chro- 
mate, soluble in ammonia and in nitric acid: 


2Agt+ CrO4— FRB AgoCrO,. 


12. Potassium Dichromate precipitates reddish-brown silver di- 
chromate soluble in ammonia and in nitric acid: 


2Ag*+Cre07- > AgeCr2O7. 


REACTIONS IN THE DRY WAY 


Fused with soda on charcoal, all silver compounds yield a white, 
malleable, metallic button without incrustation (difference from 
lead and tin), readily soluble in nitric acid (difference from tin). The 
solution is not precipitated by very dilute sulfuric acid, but is imme- 
diately precipitated by hydrochloric acid (difference from lead). 

The reactions of lead and mercurous compounds have already been 
described (pp. 201, 207). 


282 HYDROCHLORIC ACID GROUP 


Analysis of Group I 


Silver, lead and mercurous chlorides are difficultly soluble in water. +Accord- 
ing to the table on page 21, 0.088 mg. Hg.Cl, 0.015 mg. AgCl and about 1 
gm. PbCl, will dissolve in 100 cc. of water. From these values, it is evident 
that the precipitation of mercurous and silver cations will be nearly complete 
from 100 ce. of solution and that, although lead chloride is much less soluble 
in a solution containing an excess of chlorine ions than it is in water, lead 
will never be precipitated completely as chloride from aqueous solutions. By 
the addition of a large excess of alcohol, however, the precipitation of lead 
chloride can be made nearly complete, but the addition of alcohol at this stage 
of the analysis is rarely permissible except in the analysis of certain alloys 
or compounds rich in lead. 

Lead chloride is characterized by the fact that its chloride is much more 
soluble in hot water than in cold. In dissolving lead compounds in a small 
quantity of hydrochloric acid, the chloride often separates out as the solu- 
tion cools, and in such cases it is easier to remove the lead chloride at this stage 


of the analysis than to precipitate it as sulfide, dissolve in nitric acid and _ 


reprecipitate as sulfate. 

As group precipitant either hydrochloric acid or ammonium chloride can 
be used. Any other soluble chloride would accomplish the same end but would 
interfere more with the subsequent analysis. If the solution is originally 
alkaline, the addition of hydrochloric acid may cause the formation of some 
other precipitate. Thus an alkaline cyanide solution, a solution of thio salt 
in ammonium sulfide and an alkaline silicate solution will usually give precipi- 
tates upon the addition of any acid. Such precipitates often dissolve upon 
the addition of a little more acid and cause no trouble, but sometimes it is 
best to treat the precipitate as an insoluble substance by the methods 
described in Part IV. 

Bismuth, antimony and stannic chlorides on being diluted with water 
are changed into insoluble basic salts. These precipitates may be filtered off 
and dissolved by repeatedly passing a little 2-normal hydrochloric acid through 
the filter. Then, as the solution thus obtained is likely to contain lead chloride, 
it is best to evaporate to about 1 cc., dilute with 25 cc. of water and precipitate 
with hydrogen sulfide without paying any attention to the formation of a basic 
salt upon diluting: such basic salts are changed by hydrogen sulfide to less 
soluble sulfides. 


TABLE X.—ANALYSIS OF THE SILVER GROUP 





Solution may contain all the metals. Add 6-normal HCl, filter and examine the 
filtrate for succeeding groups. Treat with hot water. (1) 





Residue: AgCl, Hg:Ch. Pour ammonia | Solution: Pbt++. Test for lead with 





the filter. (3) HSO,. Filter off PbSOsand treat the 
precipitate with hot NH4C2H;O2 solu- 
Residue: Solution: [Ag(NH;).]. tion. Add KsCrO.; a yellow proces 


Hg(NH.)Cl+Hg.| Add HNO;: white tate of PbCrOs shows presence of Pb. 
precipitate shows the (2) 
presence of Ag. (4) 














SILVER 283 


The analysis of the first group of metals is based upon the solubility of 
lead chloride in hot water, the solubility of silver chloride in ammonia, and the 
blackening of mercurous chloride by ammonia. 


PROCEDURE 


1. To the cold, concentrated solution add 6-normal hydrochloric acid 
and, if a precipitate is formed, continue adding hydrochloric acid, drop by 
drop until no further precipitation takes place. Filter through a small filter 
(ef. p. 58) and wash the precipitate four times with a little cold water, blow- 
ing a fine stream from the wash bottle around the upper edge of the filter and 
waiting each time until the filter has drained before adding a fresh portion of 
water. Do not use more than 5 cc. of water each time. Take the entire 
filtrate for the analysis of Groups II, III, IV and V. Pour a little hot water 
through the filter. Test the residue by (3). 

2. Test the first 5 cc. portion of hot water that runs through the filter 
for lead by adding a few drops of concentrated sulfuric acid. If a precipitate 
forms, filter, wash once with cold water and then pour 10 cc. of hot 3-normal 
ammonium acetate solution through the filter. Add a few drops of potassium 
chromate solution to the solution and 3 cc. of 6-normal acetic acid. A yellow 
precipitate of lead chromate should form if lead is present. Continue washing 
the chlorides with hot water until 5 ec. of the washings will give no test for 
lead with potassium chromate solution. 

(3) Pour 5 ec. of 6-normal ammonia around the upper edge of the filter, 
catching it, as it runs through the funnel, in a test-tube. Acidify the filtrate 
with nitric acid. A white, curdy precipitate of silver chloride is formed when 
silver is present. The treatment with ammonia causes any mercurous chloride 
to turn black on the filter (cf. p. 201). 

If considerable mercurous chloride and very little silver chloride is present, 
the treatment with ammonia may fail to dissolve silver chloride. In such a 
case, wash the black residue and pour repeatedly a mixture of 3 cc. 6-normal 
hydrochloric acid and 10 cc. of saturated bromine water through the filter. 
This serves to convert the mercurous chloride to soluble mercuric salt and 
leaves the silver on the filter as silver chloride. Wash the filter with hot water 
and again pour 5 cc. of ammonium hydroxide through it and test the am- 
moniacal solution for silver with nitric acid. 


PART IIl es 


REACTIONS OF THE ACID CONSTITUENTS OR ANIONS 


DIVISION OF THE ACIDS INTO GROUPS 


The classification of the acids which is given here was first pub-' 
lished by R. Bunsen in 1878 (in manuscript form) for the use of his 
students; it was adopted (with his consent) by V. Meyer and F. P. 
Treadwell in their ‘‘ Tabellen zur qualitativen Analyse.’”’ It is based 
upon the different solubilities of the barium and silver salts. 


Group I 


Acids whose silver salts are insoluble in water and in nitric acid, 
but whose barium salts are soluble in water. 

To this group belong hydrochloric, hydrobromic, hydriodic, ferro- 
cyanic, ferricyanic, cobalticyanic, thiocyanic, and hypochlorous acids. 


Group II 


Acids whose silver salts are soluble in nitric acid, but are insoluble, 
or difficultly soluble, in water, and whose barium salts are soluble in water. 
To this group ‘belong hydrosulfuric, -hydroselenic, hydro 
nitrous, acetic, cyanic, and hypophosphorous acids. 


Group III 


Acids whose silver salts are white and soluble in nitric acid, but whose 
barium salts are difficultly soluble or insoluble in water, but soluble in 
nitric acid. ; 

To this group belong sulfurous, selenous, tellurous, phosphorous, 
carbonic, oxalic, iodic, boric, molybdic (selenic and telluric), tartaric, 
citric, meta- and pyrophosphoric acids. 


Group IV 


Acids whose silver salts are colored and soluble in nitric acid, but 
whose barium salts are insoluble in water and soluble in nitric acid. 
To this group belong phosphoric, arsenic, arsenious, vanadic, 
thiosulfuric, chromic, and periodic acids. 
284 





HYDROCHLORIC ACID 285 


Group V 


Acids whose silver and barium salts are soluble in water. 
To this group belong nitric, chloric, perchloric, persulfuric and the 


manganic acids. 
Group VI 


Acids whose silver salts are soluble in water, but whose barium salts 
are insoluble in nitric acid. 
To this group belong sulfuric, hydrofluoric, and fluosilicic acids. 


Group VII 


Non-volatile acids, which form soluble salts only with the alkalies. 
To this group belong silicic, tungstic, titanic, niobic, tantalic, 
and zirconic acids. 
GROUP I 


Silver Nitrate produces a precipitate insoluble in nitric acid. 
Barium Chloride causes no precipitation. 


HYDROCHLORIC ACID, HCl 


Occurrence.—Hydrochloric acid is found free in nature, but in 
small quantities (for example, in the exhalations of active volcanoes) ; 
its salts, however, are exceedingly common, especially those with the 
alkalies. (See these.) . 

Preparation.—Hydrochloric acid is set free by the action of con- 
centrated sulfuric acid upon a chloride. Ordinary rock salt is usually 
used, it being the cheapest chloride. 

If sulfuric acid is allowed to act upon ordinary salt, a consider- 
able evolution of hydrochloric acid takes place even in the cold, with 
the formation of sodium bisulfate:. 


NaCl+H2S04 = NaHS04+HClf¢ . 


And, on warming, the sodium bisulfate reacts with more sodium 
chloride: 
NaHS0O4 +NaCl = NaeS04+ HCl. 


When only one hydrogen of a dibasic acid is replaced by metal, 
the resulting salt is usually designated by placing the prefix bi before 
the name of the acid; sodium bicarbonate, NaHCO3; sodium bisul- 
fate, NaHSOa, etc. There is twice as much acid per equivalent of base 
a3 in the case of the normal salt. 

Hydrochloric acid may also be prepared by dropping concentrated 


> 


286 REACTIONS OF THE ACID CONSTITUENTS 


sulfuric acid into 12-normal hydrochloric acid. It is formed by the 
action of water on many acid chlorides: 


PCl3+3H20 = H3P03+3HCI. 


Properties.—Hydrochloric acid is a colorless gas, with a suffocating 
odor, which forms dense clouds in moist air. It is readily soluble in 
water (one volume of water dissolves, at 18° C., 451 volumes of hydro- 
chloric acid). The specific gravity of the saturated, aqueous solu- 
tion is 1.21, and 100 cc. of this solution contain 50.7 gms. of hydro- 
chloric acid gas. The concentrated hydrochloric acid of the laboratory 
is about 12-normal, and contains 36 to 38 per cent of the gas by 
weight. As regards its behavior on evaporation, see page 57. The 
aqueous solution of hydrochloric acid is one of our strongest acids. 
In dilute solution it is almost entirely ionized, 


HCl = H*+Cr, 


and such a solution is a good conductor of electricity. 
The behavior of hydrochloric acid on oxidation is extremely character- 
istic; water is formed and chlorine 1s set free: 


2HC1+0 = H20+Cle. 


Tnis oxidation will not take place on exposure to atmospheric, 
or even pure, oxygen, but only by strong oxidizing agents such as: 

The peroxides of the heavy metals, all nitrates, nitrites, chlorates, 
hypochlorites, chromates, selenates, and tellurates. 


The peroxides of the light metals do not yield chlorine, but hydrogen 
peroxide: 
BaO.+2HCl =BaCh.+H:202. 


1. Oxidation of Hydrochloric Acid by Means of Peroxides 


(a) By manganese dioxide: 
4HCl+ Mn0O2=2H20+ MnCle+Cle fT. 


The manganese is reduced from the quadrivalent to the bivalent condition 
and the negatively charged chlorine atom becomes neutral chlorine gas. Thus 
one atom of manganese oxidizes two atoms of chlorine. 


Some other acid, preferably sulfuric, can be used to combine with 
the manganese and then the yield of chlorine from a given quantity 
of hydrochloric acid will be twice as large: 


MnO2+2HCl+H2804= Mn804+2H20+Cle f. 


rl 4 


— ae PF ip 


p™ 





HYDROCHLORIC ACID 287 


(b) By lead peroxide: 
4HCI+PbO2 = 2H20+PbCle+Cle fF . 
(c) By chromium trioxide: 
2CrO3+12HCl = 6H20+2CrCl3+3Cle. 


2. Oxidation of Hydrochloric Acid by Nitric Acid, etc. 


When concentrated nitric acid acts upon concentrated hydro- 
chloric acid, the products of the reaction are water, nitrosyl chloride 


and free chlorine: 
HN0O3+3HCl=2H20+NOCl Tf +Cle ft. 


In this reaction the valence of the nitrogen has been reduced from 
five to three, and two atoms of chlorine have been oxidized to form 
an electrically neutral chlorine molecule. 

A mixture of one molecule of nitric acid with three of hydrochloric 
acid is known as 

Aqua Regia.—The acids are usually mixed, not desta ae to their 
weights, but according to their volumes. 

Aqua regia is, therefore, chlorine water, with the distinction that 
the chlorine exists in the nascent state; which explains why aqua 
regia is a much more energetic reagent than ordinary chlorine water. 
Nitrous acid, chloric acid, hypochlorous acid, selenic and telluric 
acids all react similarly with hydrochloric acid. 

Hydrochloric acid is monobasic; its salts are called chlorides. 


Solubility of Chlorides 


Nearly all chlorides are soluble in water, but the following are 
insoluble: mercurous chloride, HgeCle; silver chloride, AgCl; cuprous 
chloride, CueCle; lead chloride, PbCle; thallium chloride, TIC; 
aurous chloride, AuCl; platinous chloride, PtCle; bismuth oxychlor- 
ide, BiOCl; antimony oxychloride, SbOCl; mercuric oxychloride, 
HgeCl.O. 

All these chlorides which are insoluble in water are more soluble 
in strong hydrochloric acid. 

Aurous chloride and platinous chloride dissolve readily in aqua 
regia, as a result of being oxidized, but silver chloride is not very solu- 
ble even in aqua regia. 

By boiling the insoluble chlorides with a concentrated solution 
of sodium carbonate, all, with the exception of silver chloride, are 
readily decomposed, e.g.: 


Hg2Cle +NaeCO3= = 2NaCl+CO2 q +Hg.0. 


‘ 


288 REACTIONS OF THE ACID CONSTITUENTS _ 


By filtration.a chloride solution is obtained which is free from 
heavy metal. 
By fusing with sodium carbonate, even silver chloride is decom- 
posed, 
4AgC1+2Na2CO3 =4NaCl+2CC2 T +02 T +4Ag, 


and silver chloride may also be decomposed by nascent hydrogen 
(ef. p. 281). 

The deliquescent chlorides (lithium, calcium, and_ strontium) 
are all soluble in absolute alcohol and in amyl alcohol. 

The chlorides of potassium, sodium, and barium are quite insoluble 
in concentrated hydrochloric acid; they can, therefore, be easily 


separated from the remaining chlorides which are soluble in water 


by saturating the solution with hydrochloric acid gas. 
Almost all chlorides are insoluble in ether, with the exception of 
mercuric, stannous, stannic, auric, and ferric chlorides, 


REACTIONS OF CHLORIDES IN THE WET WAY 


A neutral solution of an alkali chloride should be used for these 
reactions. All chlorides except those of mercury and cadmium are 
almost completely ionized in aqueous solution, so that it is a matter 
of indifference which chloride we take for the following reactions, 
provided there is no independent reaction taking place due to the 
presence of the cation. 

1. Dilute Sulfuric Acid a 10) produces no reaction, even on 
warming. | 

2. Concentrated Sulfuric Acid decomposes the solid chloride 
almost completely in the cold, completely on warming. Sulfate 
and colorless hydrochloric acid gas result from this reaction; and 
the latter is easily recognized by its odor, by the clouds which it forms 
in moist air or with ammonia vapors (obtained by holding a glass rod 
wet with ammonia near the test-tube), and by its turning moist blue 
litmus-paper red. Water is not made turbid by hydrochloric acid 
(difference from fluosilicic acid). 

Silver chloride and mercurous chloride are decomposed with difficulty by 
sulfuric acid, the latter with evolution of sulfur dioxide; the mercurous sul- 


fate (which is at first formed) is oxidized (at the expense of the oxygen of the 
sulfuric acid) to mercuric sulfate: 


(a) Hg.Cl, -{- H.SO, = Hg.SO,+ 2HCl > 
(b) Hg.S0O,+ 2H.SO, = 2HgSO.+ 2H.0 + SO, T ; 


If a chloride and an oxidizing agent are heated with concentrated sulfuric 
acid, free chlorine is evolved (cf. p. 286). 


HYDROCHLORIC ACID 289 


3. Phosphoric Acid heated with a chloride similarly causes the 
evolution of hydrochloric acid gas. If an oxidizing agent is present at 
the same time, chlorine gas is evolved. 

4. Silver Nitrate produces a white, curdy precipitate of silver 
chloride, . 
| Cl-+Ag* — AgCl, 


insoluble in nitric acid, soluble in ammonia, potassium cyanide, and 
sodium thiosulfate solutions. (See Silver.) 


From a solution of silver chloride in ammonia, acids reprecipitate silver 
chloride. From a solution in potassium cyanide, acids precipitate silver cyanide. 
If it is desired to test a solution of ferrous sulfate for the presence of a small 
amount of chloride, it must be strongly acidified with nitric acid, as otherwise 
a precipitate of metallic silver will be obtained, which may cause confusion 
(ef. p. 281). The best way to test the solution of ferrous sulfate for hydro- 
chloric acid is to add sodium carbonate solution until alkaline, boil, and filter. 
In the filtrate the acids originally present are now in the form of their sodium 
salts, in the presence of an excess of sodium carbonate; and the latter should 
be neutralized with nitric acid before the silver nitrate is added. 

The detection of chlorine when present in the form of chloride of a heavy 
metal is accomplished in a similar manner; and, with the exception of silver 
chloride, any insoluble chloride may be decomposed in the same way, by 
boiling with sodium carbonate solution. 

In order to detect the presence of chlorine in silver chloride, treat it with 
dilute sulfuric acid and zine (p. 281), after a short time pour off the solution 
from the deposited silver and test it with silver nitrate. 

Or, fuse the silver chloride with sodium carbonate, extract the melt with 
water, filter, acidify with nitric acid, and test with silver nitrate. 


5. Potassium Dichromate and Sulfuric Acid.—If a dry chloride 
is mixed with potassium dichromate, concentrated sulfuric acid is 
added, and the mixture heated in a small retort, brownish vapors are 
given off which condense, in the receiver, to a brown liquid (chromyl 


chloride, CrO2Cle): 
KeCre07+4NaCl+ 3H2804 = 3H20 + 2Na28041+ Ko804+ 2CrO2Cle Ty. 


Chromy] chloride is an acid chloride, and: is, therefore, decomposed 
by water into chromic and hydrochloric acids: 


CrO2Cle+2H20 = HeCrO4+2HCl. 


On adding caustic soda or potash, an alkali chloride and a yellow 
alkali chromate are obtained. If the solution is then acidified, some 
ether and a little hydrogen peroxide added, and the liquid shaken, 
the upper ether layer will be colored blue, showing the presence of 


290 REACTIONS OF THE ACID CONSTITUENTS 


chromium; and the presence of chromium indicates that a chloride 
was originally present (difference from bromide and iodide). 
6. Potassium Permanganate oxidizes a chloride even more readily. 


On the other hand, hydrobromic and hydriodic acids are oxidized | 


much more readily than yi acid. 


Behavior of Chlorides on Ignition 
The chlorides of the alkalies and alkaline earths melt (without 


perceptible decomposition) on being heated in the air. The chlorides — 


of the sesquioxides are decomposed, more or less completely, on being 
ignited in the air. Thus, ferric chloride is almost quantitatively 
decomposed into ferric oxide, with loss of chlorine: 


2FeCls+30 = FeeQO3 +3Cle. 


In the presence of hydrochloric acid, or ammonium chloride, ferric chloride 
may be volatilized completely without any decomposition. 


The chlorides of gold and of the platinum metals are readily 
decomposed into chlorine and metal: 


2AuCls = 2Au+3Cle; 
PtCh =Pt+2Ch 1. 


The remaining chlorides are mostly volatile, without perceptible 
decomposition. 


Detection of Chlorine in Non-electrolytes 


Besides uniting with metals and with hydrogen, chlorine also 


forms compounds with the metalloids; e.g., PCls, PCls, AsCls, AsCls, — 


SbCla, SbCl5, CCla, SiCl, ete. 


All these compounds, which may be regarded as acid chlorides, — 


are decomposed by water with the formation of hydrochloric acid, 
which can be tested for in the usual way. The hydrolysis usually 


takes place at the ordinary temperatures. Thus, phosphorus tri-— 


chloride yields with water phosphorous acid and hydrochloric acid, 
PCls+3HOH = H3P03+3HCl, 


and phosphorus pentachloride yields phosphoric acid and hydrochloric 


acid: 
PCl;+4HOH = H3P04+5HCI. 


The remaining acid chlorides are decomposed in a similar way by 
water at the ordinary temperature, with the exception of carbon 


FREE CHLORINE 291 


tetrachloride, which is decomposed by water only by heating in a 


closed tube: 
st CCl4+2H20 =CO02+4HCl. 


Chlorine acts upon a great many hydrocarbons, forming substitution- 
products which are non-electrolytes, and consequently will not give the 
chloride tests; for example, if chloroform, CHCl;, is shaken with a solution 
of aqueous silver nitrate, it will not yield a precipitate of silver chloride. In 
order to test such compounds for chlorine (as is frequently necessary in the 
study of organic compounds), the chlorine must be changed to hydrochloric 
acid by one of the following methods: 

1. Carius’ Method.—By heating the compound in a sealed glass tube 
with concentrated nitric acid, in the presence of silver nitrate, the compound 
is completely decomposed; all the chlorine is changed to silver chloride, 
which can be filtered off, and, after treatment with zinc and dilute sulfuric 
acid, can be tested as above. The precautions to be taken in sealing and 
opening the tube are described in Vol. II. 

2. By Heating with Lime.—A layer of granular lime (free from chloride) 
then a mixture of the substance to be tested and lime, and finally another 
layer of lime are placed in a small glass tube, which should be about 25 em. 
long and about 1 cm. wide. 

By gently tapping the tube, a canal is opened between the upper wall of 
the tube and the substance, through which the gases evolved may escape. 
The tube is then placed horizontally in a small combustion furnace and heated 
(first the front layer of lime, then the back layer, and finally the entire contents 
of the tube) to a dull red heat. 

By this means the organic substance will be completely decomposed, and 
the chlorine will be found combined with the lime in the form of calcium chloride. 

After cooling, the contents of the tube should be dissolved in dilute nitric - 
acid, the carbon filtered off, and the filtrate tested with silver nitrate for 
chlorine ions. 

3. Treatment with Sodium.—A small amount of the dry substance to be 
tested is placed in a small test-tube, a small piece of sodium (freed from petro- 
leum) is added, and the metal covered with another layer of the substance. 

_ The tube is then heated in the gas-flame, the decomposition taking place 
suddenly with incandescence. The still hot tube is transferred to a small 
beaker containing water (which breaks the tube) and sodium chloride dis- 
solves with other sodium compounds. The solution is filtered, acidified with 
nitric acid, and then tested with silver nitrate for halogens. 


FREE CHLORINE 


Chlorine, whether produced by the oxidation of hydrochloric acid 
or by igniting certain chlorides, is a greenish-yellow gas, with a suffo- 
cating odor. It is absorbed by water (one volume of water absorbs 
at 10° C. about 2.7 volumes of chlorine gas), forming chlorine water, 
a yellowish-green liquid, and a poor conductor of electricity (although 
better than pure water, showing that some ions are present). Chlorine 


292 REACTIONS OF THE ACID CONSTITUENTS — 


decomposes water to a slight extent, forming hydrochloric acid and 
hypochlorous acid: 


H20+Cle @ 2Ht+Cl-+Clo-. 


In this reaction one atom of chlorine is oxidized to a positive valence of one 
and another atom of chlorine is reduced to a negative valence of one (cf. 
-p. 170). This reaction, as the application of the mass-action law indicates, 
is favored by the presence of hydroxyl ions and hindered by the presence 
of hydrogen ions. If, besides water, which is only slightly ionized, dilute alkali 
is present, the reaction takes place quantitatively from left to right. 


Cl,+OH- > Cl-+ClO-+H:,0. 


The decomposition of chlorine water takes place slowly in the 
dark, but more rapidly in the light, and in the presence of oxidiz- 
able substances. The bleaching action of chlorine water depends 
upon its oxidizing power. : 

Chlorine water ts a strong oxidizing agent. 

If a solution of potassium iodide is treated with chlorine water, 
iodine is set free, and the solution turns yellow to brown: / 


2KI+Cle=2KCI+Io. 


If the yellow solution is shaken with carbon disulfide, or chloroform, 
the iodine, a more highly colored, violet, non-aqueous solution of iodine 
is obtained (cf. p. 17). By the addition of more chlorine water the 
solution becomes colorless, owing to the oxidation of the iodine to 
colorless iodic acid: 

Io+6He0+5Cle = 10HCI+2HI10Os3. 


The free iodine can also be detected by the addition of some starch 
paste (instead of carbon disulfide, etc.), which is turned blue by iodine. 
1. Silver Nitrate gives a white precipitate of silver chloride when 
added to chlorine water; this precipitation, however, is not quantita- 
tive, for one-sixth of the chlorine is changed into soluble silver chlorate: 


3Clo+6Ag*+3H20 = 5AgCl+AgCl03+6H". 


On adding a slight excess of sulfurous acid to chlorine water, the chlorine 
is readily and completely changed into hydrochloric acid: 


Cl, of H.O -+ H.SO; = H.SO,+ 2HCl. 


From this solution the chlorine can be precipitated quantitatively by a 
silver solution. 

Chlorine may also be easily changed into the form of a chloride by the 
action of ammonia (cf. p. 165). 


3Cl.+8NH; oy 6NH,CI+N. t . 


FREE CHLORINE 293 


Chlorine can be changed to a chloride by the action of hydrogen peroxide 
in the presence of sodium or potassium hydroxide: ‘ 


Cl.+20H-~ — OCI-+Cl-+H20, 
and 
ClO-+H,0, — H:0+ClI-+0, fT. 


2. Metallic Mercury is attacked by chlorine at the ordinary tem- 
perature, forming insoluble mercurous chloride. 


Hge a Cle = Hge2Clo. 


If therefore chlorine water is shaken with metallic mercury until it no longer 
smells of chlorine, a neutral solution is obtained which contains no chlorine. 
If hydrochloric acid is present, the solution now reacts acid, and gives a pre- 
cipitate with silver nitrate, for metallic mercury is not attacked by hydro- 
chloric acid. ‘This reaction is used as a test for hydrochloric acid in the pres- 
ence of chlorine. 


3. Metallic Zinc also reacts with chlorine water, 


Zn+-Cl, > Zntt+2cCr-. 


HypocHLorovus AciID, HOCI 


Preparation.—A solution of free hypochlorous acid is obtained _ 


by shaking chlorine water with yellow mercuric oxide until the solution 
no longer smells of chlorine: 


2Hg0+2Clo+H20 = (HgCl)20+2HOCI. 


Brown, insoluble mercuric basic chloride is formed by the reaction, 
and the solution contains hypochlorous acid. If the solution is poured 
off from the insoluble basic mercuric salt and distilled, a pure solu- 
tion of hypochlorous acid will be obtained; which, however, cannot 
be kept long in the light, for it decomposes into hydrochloric acid 
and oxygen: 

2HOCI — 2HCI+02 7. 


Hypochlorous acid is a vigorous bleaching agent; litmus and 
indigo are quickly decolorized. 

The alkali salts of hypochlorous acid (hypochlorites) are obtained — 
by neutralizing the acid with dilute sodium or potassium hydroxide; 
or more conveniently, by the action of chlorine on a dilute caustic 
alkali solution: 

Cle +20H-=Cl-+ClO~+H20. 


The ammonium salt cannot be prepared except at very low tem- 
peratures, because the nitrogen of ammonia is so readily oxidized (ef. 
p. 165). 

All hypochlorites are readily changed, on warming, into chlorate * 
and chloride (cf. p. 170): 


3KCIO — 2KC1+KCl0s, 


consequently hypochlorites must always be prepared in cold, dilute 
solution. 
The most important commercial hypochlorite is the so-called 


“ chloride of lime,” which is obtained by passing chlorine gas over 
lime at the ordinary temperature. 





*In the presence of 40 per cent or more of caustic potash the potassium hypo- 
chlorite on being heated decomposes into chloride with evolution of oxygen and the 
formation of no perchlorate (F. WiNTELER, Z. angew. Chem., 33 (1902), p. 778). 


294 


HYPOCHLOROUS ACID 295 


REACTIONS IN THE WET WAY 


___ All hypochlorites are soluble in water, and are decomposed by 
acids (carbonic acid even). 
1. Hydrochloric Acid is oxidized by hypochlorites with evolution 
of chlorine: 
ClO-+Cl-+2Ht > H20+Cle. 


The reaction is favored by the presence of hydrogen ions. The 
reverse reaction is favored by the presence of hydroxyl ions (ef. p. 
292). 

2. Sulfuric Acid decomposes hypochlorites, setting free hypo- 
chlorous acid: | 

NaOCl+ H2804 = NaHSO4+HOCI, 


_and carbonic acid acts similarly though less energetically. 
2Na0Cl+ HeCO3 = NaeCO3+2HOCI. 


It is due to the fact that hypochlorites are so readily decomposed 
with the formation of chlorine that they act as strong bleaching 
agents, indigo solution (a solution of indigo in sulfuric acid) being 
immediately decolorized. 

Hypochlorites act as oxidizing agents not only in acid solutions, 
but also in alkaline solutions at ordinary temperatures (difference 
from chlorates), many metallic hydroxides being oxidized by them 
to higher hydroxides. Thus, ferrous hydroxide is readily oxidized 
to reddish-brown ferric hydroxide: 


2Fe(OH)2+ClO” +H20 — Cl" +2Fe(OH)s, 


and similarly lead, manganous; -nickelous, and cobaltous hydroxides 
are oxidized to brownish-black hydroxides. 
Less than 1 mg. of hypochlorite can be detected by boiling with 


lead acetate solution: 
Pb*+++ClO-+H:20 — PbO2+2H*+Cr-. 


If a peroxide is likely to be present, the solution must first be acidified 


with acetic acid before applying the test (cf. p. 86.) 
3. Iodo-Starch Paper is turned blue by hypochlorites in weakly 
alkaline solutions, owing to the separation of iodine: 


2KI+Na0Cl+ H20 =2KOH+NaCl+l.. 
21-+ClO-+H20 @ Cl-+20H +Iz. 
The extent to which the reaction takes place depends upon the 


296 REACTIONS OF THE ACID CON STITUENTS 


concentration of the hydroxyl ion; only in dilute alkali will enough or 


iodine be formed to produce the blue coloration with the starch. 


4. Metallic Mercury.—If a solution containing free hypochlorous 3 


acid is shaken with metallic mercury, brown basic mercuric chloride 
is formed, insoluble in water, but soluble in hydrochloric acid: 


2He+2HOC!I = (HgCl)20+ H20. 
Under these same conditions free chlorine forms, on being shaken 
with mercury, white mercurous chloride, which is insoluble in hydro- 
chloric acid. 


. This property is utilized in detecting hypochlorous acid in the presence 
of free chlorine. The mixture is shaken with mercury until a little of the 
solution no longer turns iodo-starch paper blue; the liquid is then carefully 
poured off, hydrochloric acid is added to the residue, and the mixture is shaken, 
when the basic chloride produced by hypochlorous acid goes into solution: 


(HgCl),0+2HCl = H,0+2H¢Ch, 


while mercurous chloride remains undissolved. 

If hydrogen sulfide is passed into the filtered solution, the formation of 
mercuric sulfide shows that hypochlorous acid was originally present. 

The salts of hypochlorous acid behave differently toward mercury than 
is the case with the free acid; they form insoluble mercuric oxide and a soluble 
chloride: 

Hg+Na0Cl =HgO+NaCl. 


5. Silver Nitrate causes in solutions of hypochlorites an incom- 


plete precipitation of silver chloride. One-third of the chlorine remains ~ 


in solution in the form of chlorate: 
3ClO~-+2Ag* — C1037 +2AgCl. 


Hypochlorous acid is distinguished from chlorine by its behavior 
toward mercury; from hydrochloric acid by its oxidizing action; 
and from chloric acid by its being partly precipitated by silver 
nitrate, and by its oxidizing action in alkaline solutions. . 


HYDROBROMIC ACID, HBr. 


Occurrence.—Bromine is always found in nature in company with 
chlorine; thus salts of hydrobromic acid are found in the ocean and 
in many. mineral waters. 

Preparation.—Hydrobromic acid is formed by the action of sul- 
furic acid upon a bromide: 


2NaBr+ H2S804 = NaeSO4+2HBr. 


But the hydrobromic acid obtained by means of this reaction is never 
pure, being contaminated with bromine; a part of the hydrobromic 
acid is oxidized by the sulfuric acid: 


2HBr+ H2SO4 — 2H20+S02 tT +Bre T : 


The more concentrated the sulfuric acid used, the larger will be 
the yield of bromine. 

If dilute sulfuric acid (83H2SO4 : 1H2O) is used, the hydrobromic 
acid obtained is nearly free from bromine. 

Pure hydrobromic acid may be obtained by: the action of an acid 
bromide upon water: 


PBr3 +3HOH = Hs3POz3 +3HBr. 


\ 


Properties. ..Hydrobromie acid (like hydrochloric acid) is a color- 
less gas with a suffocating odor, having the property of fuming in 
moist air and forming clouds of ammonium bromide with vapors 
of ammonia. It is very soluble in water. The concentrated solu- 
tion has the sp.gr. 1.78 and contains 82 per cent by weight of hydro- 
bromic acid. Hydrobromic acid is less stable than hydrochloric 
acid and more stable than hydriodic acid. 

While hydrochloric acid can be kept in aqueous solution for an 
indefinitely long time, a solution of hydrobromic acid soon turns 
brown, owing to the separation of bromine. Hydrobromic acid is 


oxidized by atmospheric oxygen: 
4HBr+O2 = 2H20+4+2Broe. 


Owing to the slight solubility of cuprous bromide, and aided by the 
presence of a little free bromine, hydrobromic acid will dissolve copper, 
297 


298 REACTIONS OF THE ACID CONSTITUENTS 


although this element is below hydrogen in the electromotive series 
(cf. pp. 38 and 41). : 
Hydrobromic acid is readily oxidized with separation of bromine, 
by peroxides, nitrates, chromates, etc., provided a concentrated solu- 
tion of hydrobromic acid is used. 
Hydrobromic acid is a monobasic acid; its salts are called bromides. 
The solubility of a bromide is similar to that of the correspond- 
ing cbloride. 


REACTIONS IN THE WET WAY 


1. Dilute Sulfuric Acid (1:10) evolves no hydrobromic acid 
from bromides in the cold, but, on warming, does so from bromides 
of the alkalies. 

2. Concentrated Sulfuric Acid causes evolution of hydrobromic 
acid and bromine from all bromides: | 


H2S804+2Br — SO47>+2HBr 
and . . 
H2$04+2Br — H20+S802 T +Bre T : 


The solution is brown, and, on warming, yellowish-brown vapors are 
given off (difference from hydrochloric acid); which, as they contain 
hydrobromic acid, fume in moist air, have a suffocating odor, and do 
not render water turbid (difference from fluosilicie acid). 

3. Phosphoric Acid causes the evolution of hydrobromic acid. 
If an oxidizing agent is also present, bromine is evolved. | 

4. Silver Nitrate produces a curdy, yellowish precipitate of silver 
bromide, insoluble in nitric acid, but soluble in ammonia, potassium 
cyanide, and sodium thiosulfate. Silver bromide is less soluble than 
silver chloride. 

If, therefore, silver chloride is digested with potassium bromide, 
the former will be changed into silver bromide: 


AgCi+KBr=AgBr+ KCl. . 


If, however, silver bromide is heated and treated with chlor 
gas, it is readily changed into silver chloride: 


AgBr+Cl=AgCl+Br. 


5. Chlorine Water, on being added to solutions of soluble bro- 
mides, sets free bromine, which is soluble in carbon disulfide or chlo- 
roform, forming a brown.solution; but it is changed, by an excess of 
chlorine water, into yellowish chloride of bromine (BrCl) (difference 
from iodine). | 


FREE BROMINE 299 


6. Potassium Dichromate, in the presence of dilute sulfuric acid, 
does not cause separation of bromine from aqueous. solutions of 
bromides; on shaking the solution with carbon disulfide, the latter 
remains colorless (difference from iodine). | 

7. Potassium Dichromate and Concentrated Sulfuric Acid.— 
On mixing a solid bromide with solid potassium dichromate, cover- 
ing the mixture with concentrated sulfuric acid and distilling, a brown 
distillate is obtained (as with a chloride), which, however, consists 
of bromine and contains no chromium: 


Cr207~+6Br-+14H* — 2Cr* **+7H20+3Bre ¢ . 


On adding dilute sodium hydroxide to the distillate, a colorless 
(or sometimes a faint yellow) solution is obtained; which, after being 
acidified with sulfuric acid, does not give the chromium reaction with 
dilute sulfuric acid and hydrogen peroxide, but merely turns brown 
owing to the liberation of free bromine (difference from chlorine). 

8. Potassium Permanganate oxidizes hydrobromic acid more 
readily than it does hydrochloric acid and less readily than it does 
hydriodic acid. 

9. Nitrous Acid does not cause the separation of bromine from 
a dilute bromide solution in the cold (difference from iodine). 


Detection of Bromine in Non-electrolytes 


The method of procedure is exactly the same as was given in the 
case of chlorine in a non-electrolyte (see p. 291). 


FREE BROMINE 


Free bromine (which may be obtained by the oxidation of hydro- 
bromic acid) is a brown liquid at the ordinary temperature, and dis- 
solves in water, forming a colored solution. The cold, saturated 
solution of bromine contains 2 to 3 per cent of dissolved bromine. 
Concentrated hydrochloric acid at the ordinary temperatures dis- 
solves much more bromine, the saturated solution containing about 
13 per cent of the substance. 

Bromine, like chlorine, acts as a strong bleaching agent (oxidiz- 
ing the color) and combines directly with metallic mercury, forming 
insoluble mercurous bromide. 

The detection of hydrobromic acid in the presence of bromine 
is accomplished in precisely the same way as was described for the 
detection of hydrochloric acid in the presence of free chlorine (cf. 


p. 293). 





HyYDRIODIC ACID, HI 


Occurrence-—Iodine occurs in nature as the iodide and as the 
iodate; most frequently as the former, accompanying (in small amounts) 
chlorine and bromine; e.g., in the ocean, in mineral waters, etc. 

_Preparation.—Hydriodic acid may be obtained pure by the action 
of an acid iodide upon water: 


PI3+3H20 = H3P03+3HI. 


If we attempt to prepare hydriodic acid by the action of sulfuric 
acid upon iodides, even from dilute solutions, it is more difficult than 
in the case of hydrobromic acid to obtain a pure product, on account 
of the strong reducing power of hydriodie acid. The hydriodic acid 
thus obtained always contains a large amount of iodine, together 
with the reduction products of sulfuric acid, the latter varying in com- 
position according to the concentration of the acid employed and of the 
iodide solution. Thus, with considerable concentrated sulfuric acid, 
sulfur dioxide is obtained: 


2Nal+3He2S04 ==> 2NaHS0O4+2H20+S02 T +Io. 


But if considerable iodide is present, the sulfuric acid is reduced 
to sulfur and finally to hydrogen sulfide. 


6KI+7H2804 — 6KHSO4+$+4H20+4+3ly, 
8KI+9H2804 — 8KHSO4+ HS 7 +4H20+4]e. 


Properties—Hydriodic acid is a colorless gas, with a suffocating 
odor; it fumes in moist air, and is readily soluble in water, forming 
a strongly fuming liquid of sp.gr. 1.99-2.10. The aqueous solution 
of hydriodic acid is even more difficult to keep than a solution of 
hydrobromic acid; it soon turns brown, owing to the separation of 
iodine: 

4HI+Oz2 (air) =2H20+2le. 

If hydrogen sulfide is conducted into the brown solution, it -is 

decolorized, with separation of sulfur: 


I2+H2S=2HI+58. 
301 


302 REACTIONS OF THE ACID CONSTITUENTS 


Owing to the extremely slight ionization of complex mercuric iodide 
anions, hydriodic acid is capable of dissolving mercury, which is con- 
siderably below hydrogen in the electromotive series (cf. pp, 10 and 


37). 
Hg+4HI =[Hgl4]H2+ He. 


Hydriodice acid (like hydrochloric and hydrobromic acids) is oxi- 
dized by peroxides, nitrates, nitrites, chromates, etc., with separa- 
tion of iodine; only in this case the oxidation of the hydriodic acid 
takes place much more readily, so that a very dilute solution is oxi- 
dized by nitrous and chromic acids even in the cold. 

Hydriodie acid is a monobasic acid; its salts are called iodides. 

The solubilities of the iodine compounds are analogous to the 
corresponding bromine and chlorine compounds. The cuprous, siiver, 
mercury and palladium salts are much less soluble than the corre- 
sponding chlorine or bromine compounds, 


REACTIONS IN THE WET WAY 


1. Dilute Sulfuric Acid (1: 10) attacks the iodides of the alkalies 
perceptibly, but only on warming. 

2. Concentrated Sulphuric Acid reacts in the cold (ef. p. 301). 

3. Silver Nitrate produces a yellow, curdy precipitate of silver 
iodide, insoluble in nitric acid, and only slightly soluble in ammonia, 
but readily soluble in potassium cyanide and sodium thiosulfate. 


Treatment with ammonia causes the silver iodide to assume a much lighter 
color. 


By the action of chlorine gas, silver iodide is readily changed into 


silver chloride: 
2AeT+Clo=2AgCl+To. 


On the other hand, if chloride or bromide of silver is treated with 
potassium iodide it will be changed into silver iodide: 


AgCl+KI=KCl-+AgI, 
AgBr+KI=KBr-+AglI. 


This apparently anomalous behavior is easy to explain. As its position in 
the electromotive series shows, free chlorine can displace bromine or iodine in 
its salts; it hasa greater tendency to be reduced to chloride ions. On the other 
hand, silver iodide is far less soluble than silver bromide and the latter is less 
soluble than silver chloride, and, for this reason ‘alone, iodine ions will replace 
chlorine or bromine ions in the case of the silver salts, 


HYDRIODIC ACID 303 


4. Lead Salts precipitate yellow lead iodide, soluble in consider- 
able hot water and forming a colorless solution which deposits gold- 
yellow plates of PbI2 on cooling. 

5. Palladous Chloride (it is best to use sodium palladous chloride) 
precipitates, from dilute solutions of an iodide, black palladous iodide 
(difference from chlorine and bromine). 


Nas(P dCl4]+2KI =2NaCl+ 2KCI +PdIo, 


which is readily soluble in an excess of potassium iodide. 
6. Cupric Salts are reduced by iodides, causing the separation 
of a brownish mixture of cuprous iodide and iodine: 


Cut te4r —> Cuzl2+ Ie. 


If sulfurous acid or sodium thiosulfated is then added to the solution, 
a nearly white deposit of cuprous iodide is obtained, owing to the 
free iodine being changed to hydriodic acid by the sulfurous acid: 


SO3° +H20+I2 —> SO4-+2Ht+2I -; 
28203 +Is — 8406 +2I-. 


7. Nitrous Acid.—If a dilute solution of an iodide is treated with 
nitrous acid, iodine separates out and the solution becomes yellow 
or brown in color (difference from chlorides and bromides): 


2HNO2+2HI =2NO JT +12+2H20. 


This extremely sensitive reaction is best performed as follows: 


Prepare a solution of nitrous acid in concentrated sulfuric acid by heating 
arsenic trioxide with nitric acid (sp.gr. 1.30-1.35) and -conducting the gases 
evolved (NO; and NO) into sulfuric acid (sp.gr. 1.75-1.80): 


As.0O;+2HNO; =As,.O;+H,0+ NO, tT +NO tT ; 


—OH 


This solution of nitrosyl sulfuric acid is sometimes called “ nitrose.” It can 
be kept for some time, but water decomposes it into nitrous and sulfuric acids: 


80:_pNot HOH ~ =H.S0.+ HNOQ.. 
Treat the solution to be tested with a few drops of nitrose and shake with 


a little carbon disulfide or chloroform; if an iodide is present, iodine is liberated, 
which colors the carbon disulfide or chloroform. 


304 REACTIONS OF THE ACID CONSTITUENTS 


8. Potassium Dichromate, in the presence of dilute sulfuric acid, 
causes the separation of iodine from dilute iodide solutions in the 
cold; the iodine can be more easily recognized by shaking the solu- 
tion with chloroform or carbon disulfide (difference from bromine): 


Cro077+61-+14Ht > 2Cr++++7H20+4+3I2: 


By heating a mixture of solid iodide and solid potassium dichro- 
mate with concentrated sulfuric acid, iodine is set free (according 
to the above equation), which distills over, and can be condensed in 
the receiver. No chromium is carried over by this reaction (differ- 
ence from chlorine). : 

9. Potassium Permanganate oxidizes hydriodic acid more readily 
than acts upon hydrobromic¢ or hydrochloric acids. The oxidation will 
take place in dilute acetic acid solutions containing considerable sodium — 
acetate. 

10. Mercuric Chloride produces scarlet mercuric iodide, soluble 
in an excess of potassium iodide: 


Hg*t+2I- — Hels, | 
Hgl2+2l" — [Hgly]-. 

11. Chlorine Water sets free iodine from iodides, 
2KI+Cle=2KCI+Iy, 


which colors carbon disulfide reddish-violet, or starch-paste blue. 
By adding an excess of chlorine water, the violet color disappears, | 
the iodine being oxidized to colorless iodie acid: 


Iz+6H20+5Cle = 10HCI+ 2HI0s. 


Detection of Iodine in Non-electrolytes 


The processes to be employed are the same as those described for 
detecting chlorine in non-electrolytes (see p. 291). 


FREE IODINE 


Free iodine forms scales resembling graphite in appearance, has 
a sp.gr. 4.94 at 17°. It melts at 114° (at the same temperature as 
sulfur), but begins to volatilize at ordinary temperatures, and is 
completely transformed into violet vapors at 200°. 

Iodine is only slightly soluble in water (100 parts of water dissolve 
0.02 part of iodine) but is soluble to a considerable extent in alcohol 
and ether, forming brown solutions, and it is much more soluble in 


FREE IODINE 305 


carbon disulfide and in chloroform; all the iodine can be removed 
from an aqueous solution by shaking a few times with either of these 
solvents (cf. p. 17). Iodine is very soluble in hydriodic acid, or in a 
solution of an alkali iodide, forming a triiodide: 


, +I2.2 I 


_ The solution of the triiodide shows all the reactions of free iodine; 
but the vapor tension of the iodine is much less than in a solution 
of pure water, because only a little free iodine is actually present at 
any one time. 

Commercial iodine always contains water, chlorine, bromine, and 
often cyanogen (iodine cyanide) as impurities. 

An aqueous soluticn of iodine is a weak oxidizing agent. 

If hydrogen sulfide is passed through an aqueous solution of iodine, 
it becomes colorless:and turbid, owing to the separation of sulfur: 


H28+1,=2HI-+S. 


Solid iodine is not acted upon at ordinary temperatures by hydrogen 
sulfide; heat is necessary to produce the endothermic hydriodie acid. 

In aqueous solution the necessary heat energy is furnished by the 
solution of the hydriodic acid formed in water. The fact that solid 
iodine is not acted upon by hydrogen sulfide, while it decomposes 
arsine, is utilized in the preparation of hydrogen sulfide containing no | 
arsenic from pyrites containing arsenic (cf. p. 234). The mixture 
of hydrogen sulfide and arseniuretted hydrogen is passed over iodine, 
and the latter only is removed. 

Sodium Thiosulfate decolorizes iodine solutions, forming sodium 
tetrathionate and sodium iodide: 


282037 +Ise — S406 +21-. 


Chlorine and bromine react in exactly the same way upon sodium thio- 
sulfate when they are not present in excess. In the latter case the reaction 
goes further and the tetrathionate is oxidized to sulfate and sulfuric acid with 
deposition of sulfur, and the sulfur itself is gradually oxidized to sulfuric acid 
by the halogens: 


28,0¢-"+12H.0+8Cl, — 16Cl-+680,-+24Ht+S,, 
S.+8H,.0+6Cl, =? 16Ht+280,-+12CI-. 


Other weak oxidizing agents, such as ferric and cupric salts, act upon 
thiosulfate similar to iodine (see Thiosulfuric acid). 


Starch Paste.—Free iodine colors starch paste blue, but only in 
the presence of hydriodic acid or a soluble iodide. 


306 REACTIONS OF THE ACID CONSTITUENTS 


Opinions differ concerning the composition of the blue “ iodide of starch.” 
Some hold that it is a compound, while others regard it as a solid solution.* 
According to Mylius,{ iodide of starch is the hydriodic acid compound of an 
iodine addition-product of starch, containing about 18 per cent iodine, corre- 
sponding to the formula [CaHO2l],-HI. This compound acts as an acid. 
If iodide of starch is produced in a neutral solution, in the presence of iodides, 
a salt of the above acid is formed, of which Mylius isolated the barium salt. 
Iodide of starch, then, can be regarded as a double salt, similar to carnallite, 
MgCl,-KCl1-6H,O. In dilute solutions it must be dissociated into its com- 
ponents, e.g., the potassium salt 


[CoH aoOvol] -KI = 4[CosH4oO201]+ KI, 


and if we assume that the compound CH, OzI is colorless, the aqueous 
solution of starch iodide will be colorless; but on increasing the concentration 
of the alkali iodide, the double salt will be less dissociated and the blue color 
of the undissociated compound will appear, which corresponds with the facts. 
If a dilute aqueous iodine solution (obtained by shaking iodine with water) 
is added drop by drop to a dilute aqueous starch solution, a blue color will 
appear at the spot where the two solutions first come in contact, but this color 
will disappear on stirring. If some potassium (or other) iodide is added to the — 
colorless solution of the starch and iodine, a permanent blue coloration will 
at once appear.{ 

The temporary appearance of the blue color, immediately on adding the 
iodine solution, is probably due to the fact that the iodine first forms a sub- 
stitution product with the starch, setting free hydriodic acid, which furnishes 
the conditions for the formation of the iodide of starch. 

The fact that a starch solution containing an iodide is much more sensitive 
than one in pure water has been known for a long time, 


Detection of Hydrochloric, Hydrobromic, and Hydriodic Acids in the 
Presence of One Another 3 


METHOD A 


The solution to be tested should contain the alkali salts of the 
above acids. Half of this solution is taken for the bromide and iodide 
test, while the other half is retained for the chloride test. 


(a) Detection of Bromine and Iodine 


1. Acidify the solution with dilute sulfuric acid, add a little carbon disul- 
fide, or chloroform, a drop of chlorine water and shake. If iodine is present 





*Kisrer, Ann. 288, 689 (1894), C. O. Harz, Chem. Zentr., 1898, I, 1018; 
AnpreEws and Gérscu, J. Am. Chem. Soc., 24, 865 (1906); Papra and Savort; Chem. 
Zentrabl., 1905, I, 1593; Karyama, Z. anorg. Chem., 56, 209 (1907). 

t Myruius, Ber., 20, 688, and C. Lonnus, Z. anal. Chem., 38, 409. 

t The blue color disappears on heating the solution, but reappears on cooling. 


FREE IODINE 307 


ae in the presence of bromine), the carbon disulfide will be colored reddish- 
violet. 

To detect bromine shake repeatedly after the addition of small portions 
of chlorine water, until the reddish-violet color has disappeared, showing that 
the iodine has been completely oxidized to iodie acid; the brown color of the 
bromine dissolved in the carbon disulfide will then appear and become a pale 
yellow on further addition of chlorine water. 

2. Instead of using chlorine water, it is often better to test for iodine 
(especially when only small amounts are present, as in mineral waters) with 
nitrous acid. Slightly acidify the solution to be tested for iodine and bromine 
with dilute sulfuric acid, add carbon disulfide and a few drops of a solution 
of nitrous acid in sulfuric acid and shake the mixture. If the carbon disulfide 
is colored reddish-violet, iodine is present. Pour off aqueous solution (through 
a moistened filter, in order to remove any suspended drops of colored carbon 
disulfide), add chlorine water and shake the solution with fresh carbon disul- 
fide. If the latter now assumes a brown color, bromine is present. 


(b) Detection of Chlorine 


The simplest way of separating chlorine from bromine and iodine is by 
fractional precipitation with silver nitrate. If the solution containing salts 
of the three halogens is treated with dilute silver nitrate, drop by drop, the iodine 
will be first precipitated as yellow silver iodide, then the bromide as a pale 
yellow silver salt, and finally the chlorine as pure white silver chloride. To 
separate the chlorine from the other two halogens, acidify a little of the solution 
to be tested with nitric acid, add a drop of dilute silver nitrate solution (1 : 100) 
shake and boil the mixture, which causes the precipitate to collect together. 
If bromine or iodine is present, the precipitate is yellow. Filter off the pre- 
cipitate and again treat with the dilute silver nitrate, etc., until a pure 
white precipitate of silver chloride is obtained, in case chlorine is present. 

By filtering off the precipitate it is easy to tell whether it is pure white, 
for the slightest tinge of yellow will show against the white paper. 


METHOD B 


This method of analysis is based upon the different degrees of readiness 
with which the iodide, bromide and chloride ions are oxidized by 1 per cent 
potassium permanganate solution. In a dilute solution containing a little 
acetic acid and considerable sodium acetate, an iodide is immediately oxidized 
with liberation of iodine, which can be removed from the solution by shaking 
with a solvent which is immiscible with water (cf. p. 17). Under these 
conditions an appreciable quantity of bromine or chlorine ions is not oxidized 
in the time required for the necessary operations with an iodide. When the 
concentration of the hydrogen ions is increased, by the addition of a prescribed 
‘quantity of sulfuric acid, the bromide is oxidized very rapidly while the rate 
of the corresponding reaction with chlorine ions is so slow that scarcely any 
free chlorine is formed. The solution may even be boiled without losing more 
than a small fraction of the chlorine ions originally present. 

The increase in the oxidizing power of permanganate upon the addition of 
hydrogen ions is a direct fulfillment of the mass-action principle (p. 13). In 


308 ~ REACTIONS OF THE ACID CONSTITUENTS 


the reaction between permanganate and halogen ions, hydrogen ions are 
required, : 
2Mn0,-+61-+8H* — 2Mn0,+4H30+38h, 
or 
2Mn0,-+101-+16H* — 2Mnt ++8H.0+5l:. 


The permanganate will be reduced only to the quadrivalent condition, and 
manganese dioxide will be precipitated, when the supply of hydrogen ions 
is limited, but bivalent manganous cations will be formed to a greater extent 
when more acid is present. 

In carrying out this method of analysis, it is very important that the direc- 
tions should be closely followed as regards the quantities of acid added. 

Procedure.—If the solution to be tested is acid, add sodium carbonate 
solution drop by drop until the solution no longer gives a decided red color to 
blue litmus paper. If too much sodium carbonate is added accidentally, add 
dilute nitric acid, drop by drop, until a very faint acid test is obtained in the 
well-stirred liquid. Then add 8 ce. of normal sodium acetate solution, 2 ce. 
of 6-normal acetic acid, and filter if any precipitate is formed. Add 3 ce. of 


chloroform and 1 per cent permanganate solution, 1 cc. at a time, until the ~ 


aqueous solution after being vigorously shaken shows the pink color of per- 
manganate. If an iodide is present, the chloroform will not be colored purple 
by the presence of free iodine. 

Pour the mixture through a moistened filter to remove the chloroform and 
precipitated manganese dioxide; the wetting with water serves to prevent 
chloroform passing through the pores of the paper. Shake the filtrate, if 
necessary, with fresh portions of chloroform to extract all of the iodine from 
the aqueous solution (cf. p. 17), filtering through a fresh filter each time. 

When all the iodine has been removed, transfer the solution to a separatory 
funnel, add 3 cc. of fresh chloroform, 5 cc. of 6-normal sulfuric acid and 1 ce. 
of the dilute permanganate, unless an excess of the last mentioned reagent 
is already present. Shake vigorously and then allow the chloroform to separate 
out beneath the water. If bromine is present, the chloroform will be colored 
yellow or orange. , 

Carefully remove the chloroform with the aid of the separatory funnel and 
- transfer the aqueous solution to a porcelain dish. Add from 5 to 20 ce. of dilute 
permanganate, according to the amount of bromide probably present, and boil 
the mixture about five minutes, or until the volume of the solution has been 
reduced to 10. cc. Filter off the manganese dioxide precipitate; and, if the 
solution is still pink, add sulfurous acid drop by drop until it is colorless. Boil- 


ing the solution serves to oxidize the last traces of bromine ions present with- 


out oxidizing chlorine ions to any extent. The excess of permanganate reacts 
with manganous ions (formed by reduction of the permanganate) to form 
manganese dioxide, 


2Mn0,-+3Mnt ++9H.0 — 5MnO:+4Ht, 


but the presence of hydrogen ions interferes with this reaction, in accordance ~ 


with the mass-action law. Consequently, either owing to the fact that there 
is a deficiency of manganous ions in the solution or because the concentration of 


the hydrogen ions is too great, it is often necessary to add sulfurous acid to 
reduce the last traces of permanganate ion, 


2Mn0,-+580;-+6H*t — 2Mnt+-+580,-+-3H.0. 


FREE IODINE 309 


Dilute the solution to 100 cc., filter if necessary, add 20 ec. of 6-normal 
nitric acid and a little silver nitrate solution. A curdy precipitate of silver 
chloride is formed if a chloride is present. 

Remark.—The above method of separation is capable of giving excellent 
results. A solid substance can be tested for chloride, bromide and iodide after 
first heating it with phosphoric acid and collecting the distillate. If about 
2 gms. of solid, 25 cc. of water and 10 cc. of 85 per cent phosphoric acid are used, 
all the iodine, bromine and chlorine will pass into the distillate. If any oxidizing 
agent is present, however, free iodine, free bromine and free chlorine will be 
formed in the order named. The free halogen can be removed by shaking the 
distillate with chloroform. Iodine and bromine are recognized by the color 
they impart to the chloroform (cf. p. 307). If free chlorine is formed, it will 
dissolve in the chloroform and liberate iodine from potassium iodide. After 
the removal of the free halogen, the distillate may be tested for hydriodie, 
hydrobromic and hydrochloric acids as described above. By distilling with 
phosphoric acid, halogen is obtained even from chlorates and bromates if a 
reducing agent is present. 


Detection of Halogen in the Presence of Cyanide 


Conduct carbon dioxide through the slightly alkaline solution until the 
escaping gas no longer gives any turbidity when passed into a nitric acid solu- 
tion of silver nitrate. 

The carbon dioxide may be prepared in a Kipp generator from marble and 
dilute hydrochloric acid and washed with sodium bicarbonate solution. The 
carbonic acid expels the weaker hydrocyanic acid from its salts (ef. p. 10). 
Owing to the very poisonous nature of hydrocyanic acid, the expulsion of the 
gas should take place under a good hood. 

After the removal of the hydrocyanic acid, test the solution in the usual 


way for halogens, 


HYDROCYANIC AcID (PRussic AciD), HCN 


Occurrence.—The compound of hydrogen with the univalent radi- 
eal cyanogen, CN, occurs to some extent in nature. I¢ is found in 
all parts of a tree growing in Java (Pangium Edule), particularly in 
the seed kernels. Its compounds are found in many plants as a 
glucoside (amygdalin), which yields, on hydrolysis, a carbohydrate, 
benzaldehyde, and prussic acid: 


CooH27NOi11+2H20 = 2C6H1206+CeHsCHO+HCN. 
Amygdalin Benzaldehyde 

This amygdalin is found in bitter almonds, in the fruit kernels 
of cherries, apricots, peaches, etc., and in the leaves of the common 
laurel tree (Prunus Laurocerasus). 

Amygdalin is usually accompanied by a ferment, so that, on macer- 
ating the parts of the plant which contain the amygdalin, an aqueous 
solution of prussic acid is obtained (bitter-almond water). 

Preparation.—If ammonia is passed over glowing carbon, ammo- 
nium cyanide is formed; so that this salt, as well as other cyanogen 
compounds, is found in the ‘“ gas-water’’ obtained by the dry dis- — 
tillation of coal. , 

Hydrocyanic acid may also be prepared by adding acid to many. 
cyanogen compounds. If yellow prussiate of potash is treated with 
dilute sulfuric acid and distilled, prussic acid is evolved, which, after 
standing over solid calcium chloride, may be obtained in the anhy- 
drous condition as a colorless, exceedingly poisonous liquid, smelling 
of bitter almonds, and boiling at 26.5° C.: 


2K4[Fe(CN)«]+6H2S04 > 6KHSO14+KeFe[Fe(CN)¢]+-GHCN. 


The best method for preparing anhydrous hydrogen cyanide * 
consists in allowing a mixture of equal volumes sulfuric acid and 
water to drop upon sticks of 98 per cent potassium cyanide. Hydro- 
gen cyanide is evolved, contaminated with traces of water which can — 
be removed by allowing the liquid to stand over solid calcium chloride. 

Properties.—The liquid (as well as the gaseous) hydrocyanic acid 





* J. Wave and L. Pantina, Proc. Chem. Soc., 190, 49 (1897-8). 
310 F 


HYDROCYANIC ACID — 311 


burns with a reddish flame, and mixes in all proportions with water, 
alcohol, and ether. 

Aqueous hydrocyanic acid cannot be kept indefinitely; a brown 
deposit soon appears, and ammonium formate is formed: 


HCN+2HOH = HCO2NHsg. 


If a little mineral acid is added to the aqueous solution, it may 
be kept much longer; but, even then, ammonia and formic acid will 
be formed after a long time. Cold concentrated hydrochloric acid 
converts hydrocyanic acid into formamide, HCN+H20=HCONH»p, 
but on warming this compound is decomposed into carbon monoxide 
and ammonia. : 

Hydrocyanic acid in aqueous solution is a very poor conductor 
of electricity; in other words, it is a very weak acid, and is ionized 
only to a slight extent. 

The salts of hydrocyanic acid, the cyanides, are very similar in 
their properties to the corresponding halogen compounds, but are 
distinguished by their ability to form stable complex salts, which 
contain scarcely any cyanogen ions in aqueous solution, and conse- 
quently do not give some of the reactions for hydrocyanic acid. 

Solubility of Cyanides.—The cyanides of the alkalies and alkaline 
earths are readily soluble in water, but hydrolysis (p. 48) takes place 
to a considerable extent: 


CN-+H20 — OH-+HCN. 


Since hydrocyanic acid is only slightly ionized, the aqueous solu- 
tion of an alkali cyanide behaves like a solution of alkali hydroxide 
containing free prussic acid; the smell of the latter can be readily 
detected. 

The remaining cyanides, with the exception of mercuric cyanide, 
are insoluble in water. 


REACTIONS IN THE WET WAY 


1. Dilute Sulfuric Acid decomposes solutions of all soluble cyan- 
ides, with the exception of mercuric cyanide, setting free hydrocyanic 
acid in the cold (recognizable by its odor). Danger! 

The insoluble cyanides are decomposed by dilute sulfuric acid 
only on warming. 

2. Concentrated Sulfuric Acid decomposes all cyanides on warm- 
ing, the complex cyanides as well as the simple ones. The metals 
are then obtained in the form of acid sulfates, the carbon of the 


312 REACTIONS OF THE ACID CONSTITUENTS 


eyanogen is changed to carbon monoxide, and the nitrogen into am- . 
monium sulfate: 
> Ni(CN)2+2H2S04+2H20 = Ni8SO4+ (NHa)2804+2CO ts 

With mercuric cyanide, besides carbon monoxide, sulfur dioxide 
and carbon dioxide are obtained; for mercuric cyanide is decomposed 
at the temperature of boiling sulfuric acid into mercury and cyanogen; 
and the former dissolves in the hot sulfuric acid, with formation of 
mercuric sulfate and evolution of sulfur dioxide: 


Hg(CN)2+6H2S804 — 2NH4HS01+ Hg( A804 '2 
+CO Tt +COz2 T +802 | +803. 


3. Silver Nitrate.—If silver nitrate is added to a solution of an 
alkali cyanide drop by drop, a precipitate is formed on the addition 
of each drop, which, however, redissolves on stirring the liquid, for 
silver cyanide is soluble in an excess of alkali cyanide: 


CN-+Agt > AgCN; 
AgCN-+4ON- > [Ag(CN)oI-. 


The complex silver cyanide ion is decomposed by further addition 
of silver nitrate, being finally completely transformed into insoluble 
silver cyanide: 


[Ag(CN)2]--+Agt > 2AgCN. 


Consequently the precipitation is complete only when an excess 
of silver nitrate is added. | 

Silver cyanide is insoluble in water and dilute nitric acid, percept- 
ibly soluble in concentrated nitric acid, and readily soluble in ammonia, 
sodium thiosulfate, and potassium cyanide. Dilute nitric acid repre- 
cipitates silver cyanide from the solution in ammonia or potassium 
cyanide. 

Concentrated hydrochloric acid decomposes silver cyanide, on 
warming, into silver chloride, with evolution of hydrocyanic acid 
(difference from silver chloride, bromide, or iodide). 

By igniting silver cyanide, there is formed cyanogen gas, metallic 
silver, and brown, difficultly volatile paracyanogen, which, on further 
heating, is completely volatilized, leaving behind pure silver: ; 


2AgCN =2Ag+-(CN)o. 


Much more certain than the silver nitrate test are the tests for 
hydrocyanic acid which depend upon the formation of Prussian blue 
or of ferric thiocyanate. 


HYDROCYANIC ACID 313 


4. Prussian Blue Reaction.—Prussian blue is formed by the action 
of ferric salts upon potassium ferrocyanide (cf. p. 150). 


3[Fe(CN).6]-~+4Fet tt > Fe4[Fe(CN) sls. 


In order, therefore, to apply this reaction to potassium cyanide, etc., it is 
necessary first to transform the cyanide into potassium ferrocyanide. This 
may be accomplished by the addition of a ferrous salt, whereby ferrous eyanide 
is first formed, which dissolves in an excess of potassium cyanide, forming 
potassium ferrocyanide: 


(a) Fe+++2CN- — Fe(CN),; 
(b) Fe(CN):+4CN- > [Fe(CN),]->. 


Potassium ferrocyanide is formed even more readily by the action of potas- 
sium cyanide upon ferrous hydroxide: 


Fe(OH).+2CN~ — Fe(CN).+20H_; 
Fe(CN),:+4CN- — [Fe(CN)«]=~. 


For the formation of the potassium ferrocyanide a little iron and considerable 
potassium cyanide are required. Consequently, to bring about the reaction, add 
a little ferrous sulfate to the alkaline solution of an alkali cyanide and boil the 
mixture. Then add a little hydrochloric acid, whereby a clear solution 
is obtained, which gives, with a little ferric chloride, the blue precipitate. If 
only traces of hydrocyanic acid are present, the solution appears green at first, 
but, after standing some time, “ flocks ” of Prussian blue will be precipitated. 

The Prussian blue reaction is used for the detection of nitrogen in organic 
substances (cf. p. 316). 

Remark. —When only a trace of cyanide is present, the Prussian blue seen 
through yellow ferric chloride appears green. If the ferric chloride solution is 
diluted largely to get rid of the yellow color, the iron is then largely in colloidal 
solution, as a result of the formation of basic salt by hydrolysis; such a solution 
reacts very slowly with ferrocyanide. A more sensitive test for traces of 
ferrocyanide is obtained by adding a saturated solution of ferrous sulfate. 
Such a solution always contains enough ferric ions to give the Prussian blue test 
with a trace of ferrocyanide ions. When more than a trace of ferrocyanide is 
present, however, it is best to test with ferric chloride solution as outlined above. 


5. The Ferric Thiocyanate Reaction.—Potassium thiocyanate pro- 
duces a red coloration with a ferric salt, owing to the formation of 
slightly ionized ferric thiocyanate (cf. p. 150): 


3CNS-+Fet + *— Fe(CNS)s, 


The cyanide, therefore, must be changed to thiocyanate in 
order to apply this reaction, which can be done by heating with sulfur, 


KCN+S=KCNS, 


314 REACTIONS OF THE ACID CONSTITUENTS 


or, better, by treatment with an alkaline polysulfide, 
KCN+(NHa4)2S2 = (NHa)25+KCNS. 


To the concentrated solution of the cyanide (in a porcelain dish) add a 
little yellow ammonium sulfide and evaporate the mixture to dryness 
on the water-bath. Add a little hydrochloric acid and a drop of ferric chloride 
solution; the characteristic blood-red coloration will be produced if only traces 
of cyanide are present. . 

It is necessary to acidify in order to destroy the (NH4)25, which would 
precipitate black Fe.S; with FeCl; and the red coloration would not appear. 


| 
6. Mercurous Nitrate produces a gray precipitate of metallic 
mercury when added to a solution of an alkali cyanide (difference 


from a chloride, bromide, or iodide): 
Hg.t ++2CN~ — Hg(CN)2+Hg. 


Behavior of Mercuric Cyanide 


Mercurie cyanide, Hg(CN)s, is a non-electrolyte, is quite soluble in water, 
in alcohol and in ether, and behaves quite differently from all other, cyanides, 
All the above-mentioned reactions, with the exception of the ferric thiocyanate 
test, fail with this compound. Mercuric cyanide solution gives no precipitate 
with silver nitrate, but a readily soluble double salt is formed, 


Similarly there is no precipitate formed with ammonia, potassium iodide, alkali. 
hydroxide, or alkali carbonate because all these reagents give, under ordinary 
conditions, precipitates of mercuric compounds which are soluble in potassium 
cyanide. Thus, for example, mercuric oxide dissolves easily in potassium 
cyanide: 

HgO+2KCN+H,0 =2KOH+He(CN)s. 


Mercuric oxide itself is fairly soluble in mercuric cyanide: 


_o/ He—CN 
HgO+Hg(CN)2 = OX He CN: 

Hydrochloric, hydrobromic and hydriodic acid decompose mercuric cyanide 
readily, but dilute sulfuric acid alone has little action upon it. In the presence 
of a soluble chloride, however, mercuric cyanide is decomposed easily by sul- 
furic acid, or even by oxalic acid or tartaric acid. If, therefore, a solution of 
mercuric cyanide is treated with common salt and dilute sulfuric, oxalic or tar- 
taric acid, then, on subjecting the mixture to distillation, a distillate is obtained 
containing hydrocyanic acid, showing all the tests characteristic of this acid. 

Mercuric cyanide in solution is acted upon by hydrogen sulfide or by alkali 
sulfides, forming mercuric sulfide and hydrocyanic acid, or one of its salts. If 
the black precipitate is removed by filtration, the ferric thiocyanate test is 
obtained without difficulty. 


HYDROCYANIC ACID 315 


Behavior of Cyanides on Ignition 


The cyanides of the alkalies and alkaline earths fuse without decomposition 
when heated out of contact with the air; heated in contact with air, they absorb 
oxygen with avidity, forming cyanates: 


2KCN +0, =2KCNO. 


Consequently the alkali cyanides are strong reducing agents (cf. p. 260). 

The cyanides of the bivalent heavy metals are decomposed on ignition, out 
of contact with the air, into nitrogen and metallic carbide; the latter often being 
further decomposed into metal and carbon: 

Fe(CN), =FeC.+N, T ; 
Pb(CN): =Pb+2C+N; Tf. 

The cyanides of the trivalent metals are unknown in the free state; those 

of the noble metals are decomposed, by igniting, into metal and dicyanogen: 
Hg(CN):=Hg+(CN): T. 
It is a characteristic property of the cyanides of the heavy metals that they 


are readily soluble in alkali cyanide solutions, forming very stable complex com- 


pounds, which are to be regarded as salts of the following acids: , 


H[R‘(CN).], H.[R™4(CN).], Hs[R™(CN).] and H.{R™(CN),]. 
The first two of the above acids are so unstable that they are decomposed, 
as soon as they are set free, into hydrocyanic acid and cyanide: 
» H[R(CN).] =HCN+RCN; 
H.[R(CN).] =2HCN+R(CN)>.. 


Consequently all cyanides which are derived from these acids evolve hydro- 
cyanic acid when treated with dilute hydrochloric or sulfuric acid in the cold. 
Such compounds are: 


K[AgCN], K[Au(CN)]., K2[Ni(CN),], Ke[Zn(CN),], Ke[Cd(CN),], ete. 


These salts must be regarded as complex compounds (p. 24), for their 
aqueous solutions contain almost no heavy metal ions; they are not precipitated 
by caustic alkali, alkali carbonate, or ammonia. From this fact it follows that 
the oxides of these metals are soluble in cyanides of the alkalies, forming the 


following complex salts: 
AgO+4CN~+H,0 — 20H~+2[Ag(CN),]-; 
Ni0+4CN~+H20 — 20H~-+[Ni(CN).]- 
ZnO+4CN~+H,20 — 20H +[Zn(CN),]-; 
Cd0+4CN-+H,20 — 20H~-+[Cd(CN),}-, 


316 REACTIONS OF THE ACID CONSTITUENTS 


Hydrogen sulfide decomposes the silver and cadmium salts without difficulty 
the zine salt slightly, and the nickel salt not at all. 

The acids of the general formula H,{R™(CN).] and H.[R™(CN),] are, in 
contrast with the above acids, quite stable in the free state, and can be usually 
obtained, without the loss of hydrocyanic acid, by acidifying a solution of one 
of their salts with cold dilute mineral acid; but, on warming the acid solution, 
hydrocyanic acid is given off. 

As typical representatives of these acids we have ferrocyanic, ferricyanic 
and cobalticyanic acids.* ‘ 

We will consider ferrocyanic and ferricyanic acids separately; but before 
doing so we will briefly describe 


DICYANOGEN 


which is obtained by heating the cyanides of the noble metals, as a colorless 
gas with a disagreeable odor; it burns with a reddish flame, and is soluble in 
water (25 parts water dissolve 100 parts of dicyanogen). The aqueous solution 
eannot be kept very long, as brownish “ flocks” separate out little by little 
(azulmie acid, C,H;N;O) and the solution then contains ammonium cyanide, 
ammonium carbonate, ammonium oxalate, and urea. 

Just as chlorine acts upon caustic alkalies, forming chloride and hypochlorite, 
so dicyanogen reacts with them, forming a cyanide and a cyanate: 


Cl.+2KOH =KCI+H:.0+ KOCI; 
(CN)2.+2KOH =KCN+H:,0+KOCN. 


On conducting hydrogen sulfide into a solution of dicyanogen, a red, 
crystalline precipitate of hydrorubianic acid is obtained: 


Cf. p. 219. 2 aoe abeadd =(CSNH2)2. 


Detection of Nitrogen in Organic Substances. (Lassaigne.) 


If a little nitrogenous, organic substance is placed with some metallic sodium 
in a small glass tube which is closed at one end, and the tube is heated till the 
glass begins to soften, then the nitrogen and some of the carbon of the organic 
substance will combine with sodium to form sodium cyanide. After heating — 
for two minutes, plunge the hot end of the glass tube in a little water, whereby 
the glass breaks and the contents of the tube are immediately wet by water. 
Filter off the alkaline solution containing the sodium cyanide from the residual 
carbon and glass splinters, add a little ferrous sulfate solution and boil. Finally 
add a few drops of ferric chloride solution and enough hydrochloric acid to 
neutralize the alkali. If nitrogen was present in the original substance, Prussian _ 
blue is now obtained. 

Remark.—Certain nitrogenous substances are decomposed with evolution of 
nitrogen before’ the temperature required for the cyanide formation is reached, f 





* Cobaltocyanic acid is extremely unstable, like mangano- and manganicyanic 
acids. Its salts evolve HCN when treated with cold, dilute, mineral acids. 
ft Grae, Ber., 17, 1178 (1884), 


DICYANOGEN 317 


and thus the nitrogen escapes the Lassaigne reaction; in other cases the experi- 
ment fails on account of the extreme volatility of the organic substance.* 

According to E. A. Kehrer,{ the Lassaigne reaction gives positive results 
in all cases if the sodium is first heated by itself and then brought in contact 
with the vapors of the organic material. The experiment is carried out in a 
small tube, drawn out at the. closed end, such as used for the arsenic test (cf. 
Fig. 21, p. 235). The substance is placed in the restricted part of the tube 
and then, in the wider part, just before the restriction is reached, a small piece 
of sodium is introduced which has been rolled between the fingers and freed from 
petroleum by touching it to filter paper. The sodium is heated until it glows 
and then, by means of another, small flame, the substance to be tested is heated 
so that it melts and the vapors just rise up to the glowing sodium but hardly 
reach through it. By taking away the small flame, the vapors can be made to 
condense and be driven toward the glowing sodium again. Otherwise, the test 
is carried out exactly as described above. 

For the detection of traces of nitrogen it is best not to add any ferric 
chloride, but to add a saturated solution of ferrous sulfate, after adding the 
acid (cf. p. 313.) 


Detection of Hydrocyanic Acid in the Presence of Halogen Acid, 
Ferrocyanic, Ferricyanic and Thiocyanic Acids 


Hydrocyanic acid is by far the weakest of all the above acids (cf. p. 10) 
and it alone is expelled from its salts by means of carbonic acid at the tempera- 
ture of boiling water. \ 

Place the solution to be tested in a small Erlenmeyer flask, add 0.6 to 1 gm. 
of NaHCO; and close the flask with a two-hole rubber stopper. Insert through 
one of the holes in the stopper a piece of glass tubing that reaches nearly to the 
bottom of the flask and serves for the introduction of the carbon dioxide gas; 
through the other hole insert a piece of tubing that reaches only to the bottom of 
the rubber stopper and serves for the escape of the gas. Pass carbon dioxide 
gas (cf. p. 309), through the liquid in the flask, gradually heat to boiling and 
conduct the escaping vapors into silver nitrate solution which has been acidified 
with nitric acid. 

If cyanide is present in the original solution, a white precipitate of silver 
cyanide is formed in the silver nitrate solution within a short time. To con- 
firm the test, decant off the solution from the precipitate and wash it a few times 
by decantation with water. Cover the precipitate with a little yellow am- 
monium sulfide, warm, and filter. Evaporate the filtrate to a small volume 
and treat with a little hydrochloric acid and a few drops of ferric chloride solu- 
tion. A red color, due to ferric thiocyanate, proves the presence of cyanide. 





* Frist, Ber., 35, 1559 (1902). 
t Ber., 35, 2523 (1902). 


FERROCYANIC ACID, Ha[Fe(CN).¢] 


Ferrocyanic acid is a white, solid substance, which is readily 
soluble in water and in alcohol, the solution soon becoming blue on 
exposure to air. The salts of this acid are much more stable than the 
acid itself, being all prepared from the potassium salt, the so-called 
yellow prussiate cf potash. This potassium salt, the most important 
ferrocyanide of commerce, is obtained by the fusion of organic sub- 
stances containing nitrogen and sulfur (blood, etc.) with potash and 
metallic iron, and by lixiviating the product of the fusion with water. 

In the melt, iron sulfide and potassium cyanide are found, which, 
on treatment with water, are changed to potassium ferrocyanide 
and potassium sulfide, 


FeS+6KCN = Ka[Fe(CN)¢]+K28, 


and, on evaporating the solution, the former salt separates out (with 
three molecules of water of crystallization) in the form of large, yellow, 
tetragonal octahedrons. 

Recently this salt has been obtained as a by-product in the manu- 
facture of illuminating-gas, Prussian blue and ammonium thiocyan- 
ate being formed from the purification of the gas. 

The following equations will give some idea of the formation of 
potassium ferrocyanide in the gas-house: 


1. Fe7(CN)is+6Ca(OH)2=4Fe(OH)3 +3Cazg|Fe(CN)6] ; 
2. Caz{Fe(CN)6]+2KCl = K2Ca[Fe(CN)¢6]+CaCh; 


Very difficultly soluble 
3. Ke2CalFe (CN) 6] oh KeCO3 = CaCO; + K4lFe (CN) 6] ‘ 


Solubility of Ferrocyanides—The ferrocyanides of the alkalies 
and alkaline earths are soluble in water; but the remaining salts 
dissolve with difficulty (if at all) in water and in cold dilute acids. 


REACTIONS IN THE WET WAY 


1. Dilute Sulfuric Acid.—The ferrocyanides are not decomposed 


by cold sulfuric acid, but break up at the boiling vemnperenyes with 
evolution of ivdniesnaiie acid: 


2Ka|Fe(CN) 6] +6H2S04 — K2Fe[Fe(CN)¢]-+6KHSO4+-6HCN. 
318 


FERROCYANIC ACID 319 


2. Concentrated Sulfuric Acid decomposes ferrocyanides com- 


pletely, on warming, with evolution of carbon monoxide, which burns 
with a blue flame: 


K4l[Fe(CN).6]+11H2SO04+6H20 = 
= FeS0O4+4KHS01+6NHsHSO4+6CO 7 . 


SO, is also liberated by this reaction, as a part of the ferrous sulfate is 
oxidized by the sulfuric acid to ferric sulfate: 


2FeSO,+280; “<> Fe:(SO,);+S0, T ° 
3. Silver Nitrate produces a white precipitate of silver ferrocyanide 
[Fe(CN)6~~]+4Agt — Ags[Fe(CN)cl, 


insoluble in dilute nitric acid and ammonia, but soluble in potassium 
cyanide solution. On treatment with concentrated nitric acid, it is 
changed to orange silver ferricyanide, and is then soluble in ammonia. 

4. Barium Chloride gives no precipitation. 

5. Ferric Salts produce a precipitate of Prussian blue in neutral 
or acid solutions (cf. p. 150). 

6. Ferrous Salts yield a light blue precipitation, which changes 
to a darker blue on exposure to the air (cf. p. 147). 

7. Cupric and Uranyl Salts produce brown precipitates. 


To detect ferrocyanic acid in an insoluble ferrocyanide, boil the latter 
with caustic alkali solution; metallic hydroxide and a ferrocyanide will be 
formed. Thus, Prussian blue yields insoluble ferric hydroxide and a soluble 
ferrocyanide: 


Fe,{Fe(CN)«];+120H- — 4Fe(OH);+3[Fe(CN),]-—. 


Filter off the insoluble hydroxide, add dilute hydrochloric acid to the filtrate 

and treat with ferric chloride, Prussian blue is again formed if a ferrocyanide is 

resent. 

Prussian blue is often used in wall-papers as a pigment. If it is desired to 
detect the presence of this compound in a wall-paper, cut about 100 cm*. of 
the paper into small pieces, boil them with caustic potash solution, filter, and 
treat the filtrate according to the method just described. In a few hours a 
distinct blue precipitate of Prussian blue will be visible in the bottom of the 
test-tube, if it was originally present. 

Some insoluble ferrocyanides do not yield the hydroxide of the metal on 
treatment with caustic alkali. Thus the brown uranyl ferrocyanide yields 
insoluble yellow potassium uranate and soluble potassium ferrocyanide (cf. 

; 188); 
: Ce zine ferrocyanide is completely soluble in caustic alkali, forming 
an alkali zincate and soluble ferrocyanide: 


Zn,[Fe’’(CN).|+8OH~ — [Fe(CN).J-~+2Zn0.-+4H,0. 


320 REACTIONS OF THE ACID CONSTITUENTS 


In order to separate the zine from the ferrocyanide, pass carbon dioxide gas 
into the solution, boil, and filter off the insoluble zine carbonate. The filtrate 
then contains potassium ferrocyanide, which can be detected as above. 

8. Lead Salts precipitate white lead ferrocyanide insoluble in 
dilute nitric acid. 


9. Thorium Nitrate added to a slightly acid solution of an alkali 


ferrocyanide produces a white precipitate, difficult to filter (differ- 
ence from ferricyanic and thiocyanic acids). 


Behavior of Ferrocyanides on Ignition 


On being ignited, the ferrocyanides yield iron carbide, cyanide, 


and nitrogen: 
Ka[Fe(CN)¢] = AKCN +FeCo+No 13 


AgalFe(CN)¢] =4AgCN-+FeCo+Np 7. 


In the latter case, the silver cyanide is further decomposed into 
metal and dicyanogen: 


2AgCN =Age+(CN)s f. 


FERRICYANIC ACID, H3{Fe(CN)o] 


Ferricyanic acid forms brown needles, readily soluble in water. 

Tts salts, the ferricyanides, are very stable, and are obtained by 
the oxidation of the corresponding ferrocyanides. The most important 
of all these salts, potassium ferricyanide (red prussiate of potash), 
K3[Fe(CN)6], is obtained by the oxidation of potassium mpi: 
with chlorine: 


2K4[Fe(CN) 6] +Cle =2KC1+2K3[Fe(CN)¢]. 


Bromine, hydrogen peroxide, etc., may be used instead of chlorine. 

‘Solubility of Ferricyanides—The ferricyanides of the alkalies 
and alkaline earths, and the ferric salt of ferricyanic acid, are soluble 
in water, but the remaining salts are insoluble even in dilute acids. 


REACTIONS IN THE WET WAY 


1. Dilute Sulfuric Acid evolves no hydrocyanie acid in the cold 
(difference from cyanides), but does so on warming with the acid. 

2. Concentrated Sulfuric Acid decomposes all ferricyanides, on 
warming, with the formation of sulfates and carbon monoxide: 


Ks3[Fe(CN)¢]+11H2S04+6H20 = 
-=FeH(SOs)2+3KHS01+6NH4HS01+6C0 1 . 
3. Silver Nitrate produces orange silver ferricyanide: 
[Fe(CN)6]=+3Ag* > Ags[Fe(CN)g], 


soluble in ammonia, but insoluble in nitric acid. 
4. Barium Chloride gives no precipitation. 
5. Ferrous Salts produce, in neutral and acid solutions, a pre- 
cipitate of Turnbull’s blue (cf. p. 147). 
6. Ferric Salts produce no precipitation, but a brown coloration. 
7. Cupric Salts yield green cupric ferricyanide: 


2 [Fe(CN)6]=+3Cut* — Cus[Fe(CN)elo. 


8. Behavior of Ferricyanides in Alkaline Solutions.—Ferricyanic 
acid is a strong oxidizing agent in alkaline solutions, being readily 
reduced to ferrocyanic acid by sulfide, iodide, sulfite, ferrous 

321 


322 REACTIONS OF THE ACID CONSTITUENTS 


hydroxide, manganous hydroxide, lead oxide, starch, cellulose (paper), 
etc.; e.g.: 


2[Fe(CN)6]" +S" — 2[Fe(CN)e~" +8; 
2[Fe(CN)o]"+2I- — 2[Fe(CN)6]~ +e; 
2{Fe(CN).6]=+S03"-+20H- — 2[Fe(CN)6]/>~ +S04-+H20; 
2{Fe(CN)¢6]=+PbO2- — 2[Fe(CN)6]” +PbO2; 
[Fe(CN)o]"+Fe(OH)2+OH™ — [Fe(CN)6]~~+Fe(OH)s. 
The ferricyanides are even reduced by ammonia, forming nitrogen: 
6[Fe(CN)6]=+8NH3 — 6[Fe(CN)¢l7~+Ne T +6NHa4". 


On account of this easy reducibility of ferricyanic acid, it is 
often difficult, sometimes impossible, to detect its presence, par- 
ticularly in an insoluble compound. If Turnbull’s blue is boiled 
with caustic potash, the residue will consist, of a mixture of ferrous 
and ferric hydroxides, and the solution will contain potassium fer- 
rocyanide. See p. 147. 

The behavior of cyanides toward suspended, yellow mercuric oxide 


is very important. Almost all cyanides, simple or complex, with the. 
exception of potassium cobalticyanide, are completely decomposed — 


by this reagent. Mercuric cyanide and an oxide of the other metal 
are formed, which, if insoluble, may be separated from the mercuric 
cyanide by filtration. Thus potassium ferrocyanide is decomposed 
by mercuric oxide as follows: 


K4[Fe(CN) 6] +3Hg0O +3H20 = Fe(OH)2+4KOH+3Hg(CN)z. 
Prussian blue as follows: | 
Fe4[Fe(CN)¢!3 +9HgO+9H20 = 8Fe(OH)2+4Fe(OH)3 +9Hg(CN)e. 
This decomposition of the cyanides by mercuric oxide is often 


used in quantitative analysis for the separation of metallic cyanides. 


Behavior of the Ferricyanides on Ignition 


\ 


The ferricyanides are decomposed into iron carbide, cyanide, 


dicyanogen, and nitrogen: 
2K3[Fe’"(CN)¢] =2FeC2+6KCN+2N2 T +(CN)e T. 


By heating a ferricyanide in a closed tube, dicyanogen therefore 
is given off, which burns with a reddish flame. 


THIOCYANIC AcID, HCNS 


Thiocyanic acid is found in small amounts, in the form of its sodium 
salt, in saliva and urine. 

The free acid is a colorless, unstable liquid, with a penetrating 
odor. It can be kept better in aqueous solution than in the anhydrous 
state, but its salts, the thiocyanates, are much more stable than the 
acid itself. The alkali salts can be prepared from the corresponding 
cyanides by heating with sulfur: 


KCN+S=KCNS. 


They may also be prepared by treating hydrocyanic acid or an 
alkali cyanide with an alkali polysulfide at ordinary temperatures, 


KCN -+(NH4)2S2 = (NH4)28+KCNS, 


or by boiling an alkali thiosulfate solution with an alkali cyanide 
(cf. p. 389): 
Na2S203-+KCN = NasSO3+KCNS. 


The easiest way to prepare ammonium thiocyanate is to allow 
a mixture of 30 cc. concentrated ammonium hydroxide, 30 ce. alcohol: 
and 7 ec. carbon disulfide, to evaporate very slowly on the water bath. 
First of all, ammonium thiocarbamate is formed 


CSe+2NH3 = NH4CS2N Hae, 
and during the evaporation this loses hydrogen sulfide: 
NH.iCS2NH2=NH4CNS+H2S. 


Solubility—Most of the thiocyanates are soluble in water; excep- 
tions are the silver, mercury, copper, and gold salts. Lead thiocyanate 
is difficulty soluble in water; on boiling with water it is decomposed. 


REACTIONS IN THE WET WAY 


1. Dilute Sulfuric Acid (double normal) causes no reaction. 

2. Moderately Concentrated Sulfuric Acid (14-normal) decom- 
poses the thiocyanates, with evolution of carbonyl sulfide, which burns 
with a blue flame: 

KCNS+2H2S04+H20 = KHSO4+ (NH4)HSO4+CoOS 7 . 
323 


324 REACTIONS OF THE ACID CONSTITUENTS 


3. Concentrated Sulfuric Acid violently decomposes thiocyanates, 
with evolution of very disagreeably smelling vapors, COS, HCOOH, 


COz, SOe, and deposition of sulfur. 
4. Silver Nitrate precipitates white, curdy, silver thiocyanate, 


CNS~+Ag* — AgCNS, 
insoluble in dilute nitric acid, soluble in ammonia. 


5. Ferric Salts produce a blood-red coloration, due to the for- 
mation of non-ionized ferric thiocyanate, 


3CNS”-+Fe — Fe(CNS8)s, 


very soluble in ether (cf. p. 150). 
6. Mercuric Nitrate precipitates white mercuric thiocyanate, 


Hg*+++2CNS8-+Hg(CNS)>, 


very difficultly soluble in water, but readily soluble in an excess of 
potassium thiocyanate: 


Hg(CNS)2+CNS” — [Hg(CNS)s3]-. 


If dry K[Hg(CNS)s] is heated, the salt expands greatly (Pharaoh’s . 
serpents). 
7. Mercuric Chloride gives a precipitate only after long standing. — 
8. Mercurous Nitrate produces a gray to black precipitate. On 
adding mercurous nitrate drop by drop to a fairly concentrated solu- 
tion of potassium thiocyanate, a gray precipitate of metallic mercury 
is first obtained, and the solution contains potassium mercuric thio- 


cyanate: 
Hgo**++3CNS~ — [Hg(CNS)3]-+Hg. 


If the addition of mercurous nitrate is continued until no more mer- 
cury is precipitated, and the solution then filtered, the filtrate will 
contain potassium mercuric thiocyanate; but, on adding still more 
mercurous nitrate, pure white, mercurous thiocyanate is precipitated: 


2{Hg(CNS)s3]-+3Hg2* * — 2Hg*++3Hg2(CNS)s. 


If, on the other hand, a very dilute solution of potassium thio- 
cyanate is added to a very dilute solution of mercurous nitrate, the 
white precipitate of mercurous thiocyanate is obtained directly: _ 


Hg2t*+2CNS~ — Hg2(CNS)p. 


9. Cupric Salts.—On adding a few drops of a solution containing 
a cupric salt to one of an alkali thiocyanate, the solution is colored 


THIOCYANIC ACID 325 


emerald-green; and, on further addition of the copper solution, black 
cupric thiocyanate is precipitated. If sulfurous acid is added, 
white cuprous thiocyanate is deposited, 


2Cut++S03"+2CNS-+H:0 — Cu2(CNS)2+2H*+80¢, 


insoluble in dilute hydrochloric and sulfuric acids. 

10. Cobalt Salts.—If a solution containing an alkali thiocyanate 
is treated with a small amount of a cobalt salt, and the solution 
shaken with a mixture of equal parts amyl alcohol and ether,* the 
upper layer of alcohol ether separates out azure-blue in color (ef. 
p. 182). This reaction is. analogous to that of cyanic acid upon cobalt 
salts (cf. p. 344). 


Detection of Thiocyanates in the Presence of Halogen and Cyanide 


| First free the solution from hydrocyaniec acid by adding a little sodium 
bicarbonate, heating to boiling and passing carbon dioxide gas through the 
solution until the escaping gas led into slightly acid silver nitrate solution 
gives no turbidity of silver cyanide. 

After the removal of the hydrocyanic acid, test for thiocyanate, in the 
absence of iodide, by acidifying with hydrochloric acid and adding a few drops 
of ferric chloride. A blood-red coloration shows the presence of thiocyanate. The 
test cannot be obtained satisfactorily in the presence of an iodide because it 
also will react with ferric chloride, causing liberation of free iodine. 

If an iodide is present, add a little nitric acid to the solution from which 
the hydrocyanic acid has been expelled, and precipitate the halogens and 
thiocyanate by the addition of an excess of silver nitrate solution. After the 
precipitate has settled, decant off the supernatant solution and wash the 
precipitate several times by decantation with water. Then shake the precipi- 
tate vigorously with 6-normal ammonia solution. This dissolves the chloride 
and thiocyanate readily, all or a part of the bromide, but no appreciable quantity 
of silver iodide. Filter and add colorless ammonium sulfide to the filtrate. 
Filter off the silver sulfide precipitate, add a drop of sodium carbonate solution, 
evaporate to small volume, acidify with hydrochloric acid, and test for thio- 
cyanate with ferric chloride solution. 


Detection of Halogens in the Presence of Thiocyanate 
VoLHARD’s MrEetTHop 


Treat the nitric acid solution with an excess of silver solution, filter and 
dry the precipitate with suction. Transfer the precipitate to a porcelain dish 
and heat it on the water bath for 45 minutes with concentrated nitric acid. 
This causes the complete decomposition of all the thiocyanate, 


GAgCNS+4H,0+16HNO; =3Ag,S0,+3(NH,).S0.,+5CO, T +16NO 1, 





* Or with amyl alcohol alone. 


’ 


326 REACTIONS OF THE ACID CONSTITUENTS 


and does not attack the silver halides appreciably. Dilute, filter and wash a 
few times with hot water. Reduce the precipitate with zinc and dilute sulfuric 
acid (p. 281) and test the filtrate according to p. 307. 

Or, instead of treating the washed silver precipitate with nitric acid it may 
be boiled with sulfuric acid (1 : 1) until the precipitate becomes black and col- 
lects in a ball: 


2AgCNS+2H.80,+3H.0 =2NH,HSO.+COS Tf +CO: T +AgS. 


Dilute the solution with water, filter off the silver halides and silver sulfide, 
wash with water, reduce with zinc and sulfuric acid and test for the halogens — 
after boiling off the hydrogen sulfide from the last filtrate. 

Remark.—It is always necessary to destroy the thiocyanate before reducing 
with zinc and sulfuric acid, because otherwise hydrocyanic acid will be formed: 


2AeSCN-+3H, > 2Ag-+2HCN-+2H,S 1. 


According to A. W. Hofmann, free thiocyanic acid on being reduced with 
nascent hydrogen yields a mixture of thioformaldehyde, methylamine, ammonia 
and hydrogen sulfide. 


Testing Commercial Alkali Thiocyanate for Chloride 
C. Mann’s MetHop 


Dissolve 5 gms. of the alkali thiocyanate in 20 cc. of water and to the solu- 
tion add 20 gms. of crystallized copper sulfate dissolved in 100 cc. of water. 
A black precipitate of cupric thiocyanate is formed: 


Cutt++2CNS- — Cu(CNS), 


Pass hydrogen sulfide gas through the solution until the precipitate becomes 
nearly white: 


2Cu(CNS),-+H.$ =Cu(CNS).+8+2HCNS. 


‘Then, when the supernatant blue copper solution begins to get brown, 
owing to the formation of copper sulfide, stop introducing the hydrogen sulfide 
gas and allow the liquid to stand a few hours. During this time the thiocyanic 


acid formed by the above reaction is acted upon by the copper sulfide present, 
as follows: 


2Cu(CNS).+2CuS — 2Cu,(CNS).+28. 


Filter and treat the filtrate with silver nitrate. A white precipitate shows 
chloride to be present. 

Remark.—An equally good method is the following: Treat the solution of 
the thiocyanate with an excess of copper sulfate solution and-introduce sulfur 
dioxide gas until the precipitate becomes white cuprous thiocyanate. Allow 
the solution to stand several hours, then filter off the cuprous thiocyanate, treat 
the filtrate with nitric acid and test for chlorine with silver nitrate. 


Instead of sulfurous acid, hydroxylamine sulfate may be used to reduce the 
cupric solution. 


THIOCYANIC ACID 327 


Detection of Ferro- and Ferricyanides in the Presence of Thiocyanate 
Mertuop or P. E. Brownine anp H. E. Parmer 


Acidify the dilute solution of the alkali salts of these acids with acetic — 
acid or hydrochloric acid, avoiding an excess, and add a solution of thorium 
nitrate; finely divided thorium ferrocyanide will be precipitated. Shake the 
solution with finely divided asbestos, or filter paper pulp, filter and wash the 
precipitate with cold water. Pour dilute sodium hydroxide solution over the 
precipitate, acidify the solution thus obtained and add a few drops of ferric 
chloride; Prussian blue is formed if a ferrocyanide was present. 

Add cadmium sulfate solution to the filtrate from the thorium ferrocyanide 
precipitate, shake with finely divided asbestos and filter off the cadmium ferri- 
cyanide. Wash the precipitate with cold water, dissolve in sodium hydroxide 
solution, acidify the solution with hydrochloric acid and treat with ferrous sul- 
fate solution. The formation of Turnbull’s blue shows the presence of ferricyanic 
acid. 

Adding ferric chloride to the filtrate from the cadmium ferricyanide pre- 
cipitate, a blood-red coloration will be obtained if a thiocyanate is present. 


Behavior of Thiocyanates on Ignition 


The thiocyanates of the alkalies melt readily, and are colored successively 
yellow, brown, green, and finally blue, becoming white again on cooling. 

The thiocyanates of the heavy metals are decomposed into sulfide, splitting 
off carbon disulfide, dicyanogen, and nitrogen. Thus cuprous thiocyanate 
is decomposed according to the following equation: 


4Cu,(CNS), =4Cu,S+2C8, f+3(CN). T +N: T. 


The mercuric thiocyanates swell tremendously on being heated (Pharaoh’s 
serpents), 


COBALTICYANIC ACID, H;[Co(CN).¢] 


The free acid can be obtained by suspending the lead or copper salt in 
water and saturating the water with hydrogen sulfide gas; the lead or copper 
cobalticyanide is changed into less soluble sulfide and, by filtering, an aqueous 
solution of cobalticyanic acid is obtained from which the solid acid deposits 
upon evaporation. The free acid may be prepared also by treating the potas- 
sium salt with nitric acid, evaporating to dryness on the water bath and extract- 
ing the acid with alcohol. After evaporating off the alcohol, needle-shaped, 
hygroscopic crystals of H;[Co(CN).] are obtained. 

Cobalticyanic acid is extremely stable. It is not decomposed by boiling 
with concentrated hydrochloric or nitric acid, by chlorine or by boiling with 
HgO. By heating with concentrated sulfuric acid, however, it is decomposed 
with evolution of carbon monoxide and carbon dioxide: 


2H;[Co(CN),]+8H2SO0.+ 13H20 =2CoSO.,+6(NH4)2S80,+11C0 Tf +COz T. 


Solubility of Cobalticyanides—The alkali, alkaline-earth, ferric and other 
trivalent metal salts are soluble in water. Most of the other salts witl: 
bivalent metals and with the heavy metals are insoluble in water and in acids 


REACTIONS IN THE WET WAY 


1. Dilute Sulfuric Acid. No reaction. } 

2. Concentrated Sulfuric Acid decomposes all of the salts with evolution 
of CO and CO, and formation of blue, anhydrous cobaltous salt. 

3. Silver Nitrate produces a white precipitate insoluble in nitric acid but 
soluble in ammonia. 

4. Lead Acetate, Ferric Chloride and Mercuric Chloride produce no 
precipitates. , 

5. Cobaltous Nitrate produces a pink precipitate insoluble in nitric acid 
but soluble in ammonia. 

6. Nickel Sulfate gives a blue precipitate insoluble in nitric acid but soluble 
in ammonia. 

7. Copper Sulfate produces a light-blue precipitate insoluble in nitric acid 
but forming a blue solution with ammonia. 

8. Cadmium Sulfate and Zinc Sulfate give white precipitates insoluble 
in nitric acid but soluble in ammonia. 

9. Ferrous Sulfate produces a white precipitate insoluble in nitric acid. 

10. Mercuric Nitrate produces a white, voluminous pisenitars insoluble 
in nitric acid. 

REACTIONS IN THE DRY WAY 


All the salts of eobalticyanic acid are decomposed by ignition, leaving 
cyanide and cobalt carbide behind. The alkali and alkaline-earth salts color 
the borax bead blue in both the oxidizing and reducing flames. 


328 


GROUP II 


Silver Nitrate produces a precipitate soluble in nitric acid. 
Barium Chloride causes no precipitation. 


NITROUS ACID, HNO: 


Occurrence.—Nitrous acid never occurs free in nature except in 
the form of its salts, the nitrites. It is found in the air, as ammo- 
nium nitrite and in many soils and waters, particularly in those which 
are contaminated with ammonia or decaying substances. 

Ammonia is oxidized by the action of microérganisms (monas 
-nitrificans) to nitrous acid, which combines with more ammonia to 
form ammonium nitrite. 

Preparation of Nitrous Acid and its Salts.—Nitrous acid is formed 
by the gentle reduction of nitric acid. If zine is allowed to act upon 
dilute nitric acid for a short time, the latter is reduced to nitrous acid, 


HNO3+He2=H20+HNOsz, 


but the reduction can easily go a little farther, forming NO, N20, 
and Ne; while by long-contained action of the zinc, hydroxylamine, 
NH20H, and even ammonia are formed. 

If nitric acid of sp.gr. 1.3 is heated with arsenious acid, starch, 
etc., a mixture of nitric oxide and nitrogen peroxide is obtained, which, 
on cooling to —21° C., condenses to a bluish-green- liquid, N2Oz, the 
anhydride of nitrous acid. 

If the anhydride is treated with ice-cold water, a bluish-green 
liquid is obtained, which contains nitrous acid, but always in com- 
pany with nitric acid; for N2O3 unites with water, forming nitrous 
and nitric acids, and nitric oxide: 


2N203+H20 =HNOs+HNO2+2NO0 fT. 


At a higher temperature nitrous acid is gradually changed into 
nitric acid: 


3HNO2=HNO3+2NO Tf +H20. 


Pure nitrous acid, therefore, is not known. 
If the above mixture of nitric oxide and nitrogen per oxide is con- 
329 


330 REACTIONS OF THE ACID CONSTITUENTS 


ducted into concentrated sulfuric acid, the two gases are readily 
absorbed, forming nitrosyl sulfuric acid: 


2H2804+NO+NO2 — H20+2H(NO)SOs. 


This solution of nitrosyl sulfuric acid in sulfuric acid is sometimes 
called ‘‘nitrose.”’ 
If the solution is added to cold water, sulfuric and nitrous acids 


are formed: 
H(NO)SO4+H20 — HNO2+H2S04. 


A solution of nitrosyl sulfuric acid can be kept indefinitely, so 
that it is a convenient reagent for the immediate production of nitrous 
acid at any time. 

The salts of nitrous acid, the nitrites, are much more stable than 
the free acid, and may be obtained by the ignition of nitrates: 


2NaNOz = 2NaNOoe +02 T = 


Nitrites prepared in this way always contain some oxide and some: 
nitrate as impurity.* In order to obtain a pure nitrite, silver nitrite 
is treated with the calculated amount of a metallic chloride: 


AgNOo+NaCl=AgCl+NaNOsg. 


The soluble nitrite can be separated from the insoluble silver 
chloride by filtration. 7 

Solubility of Nitrites—AlIl nitrites are soluble in water; but silver 
nitrite and potassium cobaltic nitrite are difficultly soluble. 


REACTIONS IN THE WET WAY 


As all nitrites are soluble in water, the reactions which serve for 
the detection of this acid cannot be those of precipitation, but rather 
those in which a change of color takes place, owing to an oxidation or 
reduction. 

Nitrous acid sometimes acts as an oxidizing agent, and sometimes 
as a reducing agent. 

1. Dilute Sulfuric Acid decomposes all nitrites in the cold, setting 
free brown vapors: 


(a) NaNOo+He2S04 = NaHSO4+ HNO>:; 
(b) 3HNOe=HNOs+2NO 7 +H20; 
(c) 2NO+02(air) =2NO2 T . 





*If the nitrate is heated with a metal, e.g., lead, the reduction takes place at 
a lower temperature and is almost quantitative. 


\ 


NITROUS ACID 331 


2. Concentrated Sulfuric Acid reacts exactly the same, but much 
more violently. 

3. Silver Nitrate precipitates from nitrite solutions crystals of 
silver nitrite in the form of needles, which are slightly soluble in cold 
water (300 cc. of water dissolve 1 gm. of silver nitrite at, the ordinary 
temperature). In boiling water, silver nitrite is considerably more 
soluble. 

4. Cobalt Salts produce (with an excess of potassium nitrite and 
acetic acid) a yellow crystalline precipitate of potassium cobaltic 
nitrite (cf. p. 182). 

5. Indigo Solution is completely decolorized by warming with 
nitrous acid. 

6. Hydriodic Acid is oxidized by nitrous acid with separation of 
iodine: 

2I-+2N0.2-+4Ht — 2H20+2NO tT +To. 

If, therefore, a nitrite is added to a solution of potassium iodide 
and the solution is acidified with sulfuric or acetic acid, the solution 
becomes yellow, owing to the separation of iodine. If the solution 
is now shaken with chloroform or carbon disulfide, the latter will 
be colored reddish-violet. Or, if a little starch paste is added, it will 
be colored blue by the iodine. 

As the above equation shows, hydrogen ions are required in the reaction. 
If considerable alkali acetate is present, there is no separation of iodine on the 
addition of acetic acid, but if a few drops of a strong mineral acid are added, 
- jodine is at once set free. This is a good illustration of the mass-action principle 
(p. 18) and common ion effect (p. 45). 

This exceedingly delicate reaction is also caused by the action of a great 
many other oxidizing agents; and it can only be used for the detection of nitrous 
acid when it is known that all such oxidizing substances are absent. 

As ferric salts also cause liberation of iodine (cf. p. 37) it is evident that 
nitrous acid cannot be tested for by the above test in the presence of a ferric 
salt. If, however, a large excess of sodium phosphate is added to the solution, 
together with a little potassium iodide and some sulfuric acid, no iodine will be 
liberated by the ferric salt and the presence of a trace of nitrous acid may 
be detected. The ferric ions are converted into very slightly ionized ferric 
acid phosphate in this test and the concentration of the ferric ions is made so 
small that there is no appreciable reaction with the iodide ions. Artmann,* 
who suggested this procedure, adds 8 gms. NazHPQ,-12H,0, 0.2 gm. KI, 5 
ec. 4-normal H:SO, and a little starch solution to 100 cc. of a solution of 
the ferric salt which is to be tested for nitrite. If as much as 0.3 mg. of NOs 
is present, the intense blue color will be obtained immediately. 


7. Ferrous Salts are oxidized to ferric salts, with evolution of 
nitric oxide: 


Fett++N02-+2Ht > Fe**+*+N0O fT +H20. 
* Chem., Zlg., 1918, 501. 





332 REACTIONS OF THE ACID CONSTITUENTS 


The nitric oxide dissolves, in the cold, in the excess of ferrous salt, 
forming a brown compound of a varying composition: (FeSO4)2(NO)p. 


To obtain this compound, add a little acid to a concentrated solution of 
ferrous sulfate and carefully pour the solution to be tested on top. At the zone 
of contact between the two solutions the dark-brown coloration will be apparent. 

Nitric acid gives the same reaction, but only on the addition of concentrated 
sulfuric acid. 


8. Potassium Permanganate.—If nitrous acid is added to a warm 
acid solution (about 40° C.) of potassium permanganate, the latter 
will become decolorized, owing to the oxidation of the former to nitric 
acid: 

2Mn04-+5NO2-+6Ht — 2Mn*t+5N0O3_+3H20. 


In this reaction nitrous acid acts as a reducing agent. 


9. Detection of Small Amounts of Nitrous Acid by the Peter Griess 
Method.*—To detect the small amounts of nitrous acid which may be present 
in drinking-water, of the above reactions only that of potassium iodide and 
starch is delicate enough. But as hydrogen peroxide and ferric salts are also 
likely to be present, both of which cause the separation of iodine from an acid 
solution of potassium iodide, it is evident that dependence upon this reaction 
alone would often lead to error. 

Consequently, in order to detect the presence of traces of nitrous acid we 
make use of a reaction which was first proposed by Peter Griess, and which is 
caused by nitrous acid only. It depends upon the formation of an intensely 
colored azo dyestuff. 

Peter Griess used as his reagent phenylene diamine, whereby a yellow dye, 
Bismarck brown, is formed. Recently, according to the suggestion of Ilosvay 
v. Ilosva,} an acetic acid solution of sulfanilic acid and of a-naphthylamine is 
used instead. According to Lunge,t it is best to mix the solutions of the last 
two reagents. 

The reagent is prepared as follows: 

1. Dissolve 0.5 gm. of sulfanilic acid in 150 ce. of 2-normal acetic acid. 

2. Boil 0.2 gm. of solid a-naphthylamine with 20 cc. of water, pour off the 
colorless solution from the bluish-violet residue, and add to the colorless solu- 
tion 150 cc. of 2-normal acetic acid. 

Mix the two solutions. The mixture keeps well only when kept in a dark 
place. It turns reddish if exposed to the light, and cannot then be decolorized 
by shaking with zinc. 

Procedure.—Treat about 50 cc. of the water with 2 cc. of the above reagent, 


and allow it to stand five or ten minutes; it will be colored a distinct red if a 
trace of nitrous acid is present. 


10. Diphenylamine, dissolved in concentrated sulfuric acid, is 





* Ber., 12, (1879) 427. 
¢ Bull. chim. [3] 2, 317. 
¢ Z. angew. Chemie, 1889, Heft 23. 


NITROUS ACID 333 


colored intensely blue by nitrous acid. Nitric acid and many other 
oxidizing substances, such as selenic acid, chloric acid, ferric chloride, 
etc., will give the same reaction (cf. Nitric Acid). 
11. Brucine dissolved in concentrated sulfuric acid (according 
to G. Lunge and A. Lwoff *) gives no reddish coloration when treated 
with nitrosyl sulfuric acid. 


Dry, repeatedly-recrystallized silver nitrite, containing 70.05 per cent. 
silver (theory 70.09), did give with brucine (cf. p. 395), in an atmosphere of 
carbon dioxide, a weak but nevertheless distinct test for nitric acid. probably 
due to the presence of traces of nitrate remaining in the silver nitrite. On 
dissolving 15 mgm. of this same nitrite in water, adding an equivalent amount 
of sodium chloride and diluting to one liter, a solution of sodium nitrite was 
obtained, of which one ce. added drop by drop with constant stirring to about 
4 cc. of concentrated sulfuric acid yielded a solution of nitrosyl sulfuric acid 
which gave no sign of a red coloration with a drop of brucine reagent. The test 
was immediately obtained, however, on adding a trace of nitric acid to this 
solution. 
Brucine, therefore, is a reagent by which nitric acid can be detected in the 
presence of nitrous acid. 


12. Urea acts with nitrous acid with evolution of nitrogen and 
carbon dioxide: 


CO(NH2)2+2HNO2 > 3H20+C02 T +2N2 7. 


In this reaction the nitrogen of the urea, with its negative valence 
of three, is oxidized by the nitrogen of nitrous acid, which has a positive 
valence of three. In alkaline solutions free halogens will react with 
urea, but in acid solutions the reaction with nitrous acid is sensitive 
and characteristic. ; 

13. Ammonium chloride, on being added to a boiling solution of a 
nitrite in dilute acetic acid, causes evolution of nitrogen gas (cf. p. 99): 


NH,t+N0O2°— 2H20+Noe2 T ; 


HYDROGEN SULFIDE (HYDROSULFURIC ACID), H2S 


Occurrence and Preparation.—Hydrogen sulfide is found in volcanic 
regions, in many mineral waters (the so-called ‘ sulfur” waters), 
and, in general, wherever substances containing sulfur are subject to 
decay; or when they come in contact with decaying substances. Sul- 
fates are easily changed into sulfides by the action of microérganisms 
which are present in the air; and this is the reason why many mineral 
waters containing sulfates smell of hydrogen sulfide after standing some 
time in a corked flask. If, however, the flask and the cork are sterilized, 
the water can be kept indefinitely. The formation of hydrogen sul- 
fide from sulfates takes place as follows: 

By means of carbonaceous matter (dust, etc.) the sulfates are 
reduced with the aid of microérganisms, at first to sulfides. 


Na2SO4+Ce2=2CO02+Naz8S, 
which are then decomposed by carbonic acid: 
NaeS a HeCO3 = NaeCO3 +H.S. 


Just as hydrogen sulfide may be made from sodium sulfate by 
the action of organic matter in a corked flask, so in nature the same 
process brings about the presence of hydrogen sulfide in many mineral 
waters. 

For laboratory purposes, hydrogen. sulfide is. similarly prepared 
by the action of dilute sulfuric or hydrochloric acid upon a sulfide 
(usually iron sulfide, FeS, on account of its cheapness and stability). 

Properties—Hydrogen sulfide is a colorless gas, with an odor 
like that of rotten eggs; it is absorbed by water at the ordinary tem- 
peratures (one volume water absorbs two to.three times its own 
volume). The saturated solution at 25° is approximately 0.1.molal= 
0.2 normal. The higher the temperature, the less the solubility, as 
with all gases. For the primary ionization: H2S @H++HS-, the 
value of the ionization constant 





(2g bod: bp z 
(Hs) =k, is 0.911077. 
For the secondary ionization, HS~ - H++S-, the value of the constant 
[H+] x[S=] 





[HS] =ke is 1.210-5, 
334 


HYDROGEN SULFIDE 335 


In the saturated solution at 25°, the concentration of the hydrogen 
ion is about 0.9X10~ molar equivalents per liter and of the simple 
sulfide ion it is only 1.2X10°'. According to the table on p. 56, 
therefore, a saturated solution of hydrogen sulfide should react acid 
to phenolphthalein and to azolitmin but not to methyl orange. 

The solution of hydrogen sulfide becomes turbid on standing in 
the air as a result of its oxidation by atmospheric oxygen: 


2H28 + O02=2H20+S8o. 


Hydrogen sulfide burns in the air with a bluish flame to water 
and sulfur dioxide: 
2H2S +302 =2H20+2S02. 


The salts of hydrosulfuric acid are called sulfides. 

Solubility cf Sulfides—The sulfides of the alkalies and the hydro- 
and polysulfides of the alkaline earths are soluble in water. The 
monosulfides of the alkaline earths, particularly calcium sulfide, CaS, 
are difficultly soluble in water, but they are gradually changed from 
contact with water into soluble hydrosulfides: 


2CaS+2H20 — Ca(OH)2+Ca(SH)o. 


The remaining sulfides are insoluble in water. Of these latter 
FeS, MnS, and ZnS are decomposed by dilute hydrochloric acid with 
evolution of hydrogen sulfide; others require concentrated hydro- 
chloric acid, e.g., SbeS3, SnSe, PbS, NiS, CoS, CdS; while the remain- 
ing are insoluble in concentrated hydrochloric acid, but are all soluble 
in aqua regia with separation of sulfur. 


REACTIONS IN THE WET WAY 


Free hydrogen sulfide, as has been stated, is a very weak acid, being even 
weaker than carbonic acid. The soluble neutral salts R.S are decomposed into 
metal and sulfur ions, 

RS = 2Rt+S-, 
but under the influence of water, some of the bivalent sulfur ions are changed 
to univalent HS ions, 
HOH+S~- — HS-+0H; 
some of the bivalent sulfur ions remain in solution, and, in fact, more in con- 
centrated solutions than in dilute ones. 

As, therefore, an aqueous solution of a sulfide contains both S ions and SH 
ions, while the solutions of the free acid contain chiefly SH ions, it is plain why 
in many reactions the former react in a somewhat different way from the latter. 


1. Dilute Sulfuric Acid decomposes all soluble, and some insoluble, 
sulfides, with evolution of hydrogen sulfide. 


336 REACTIONS OF THE ACID CONSTITUENTS 


2. Concentrated Sulfuric Acid decomposes all sulfides, on warm- 
ing, with evolution of sulfur dioxide and deposition of sulfur: 


NaoS+2H2804 = Na2gSO4+2H20+802 fT +8. 


But even the sulfur goes over into sulfur dioxide after being heated 
with the sulfuric acid for some time: 


2H2S04+S =2H20+3S802 fF. 


8. Silver Nitrate produces, in solutions of hydrogen sulfide and 
of soluble sulfides, a black precipitate of silver sulfide, 


2A0++HS > AgsS+2H*, 


insoluble in cold nitric acid, in which, however, it dissolves on warm- 
ing (ef. p. 181). 

4. Barium Chloride causes no precipitation. 

5. Lead Salts (best a solution containing an excess of alkali) pro- 
duce a black precipitate of lead sulfide. . All sulfides which are decom- 
posed by hydrochloric acid evolve hydrogen sulfide, which, on coming 
in contact with a piece of filter-paper moistened with an alkaline lead 
solution, colors the latter black. An insoluble sulfide (pyrite, arsenic 
sulfide, mercuric sulfide, etc.), evolves hydrogen sulfide with hydro- 
chloric acid and nascent hydrogen. 


“\ 


To test an insoluble sulfide, such as the mineral pyrite, place a little finely 
granulated tin in a test tube, cover it with 6-normal hydrochloric acid and heat 
gently. Hold a piece of filter paper which has been moistened with lead acetate 
solution and a little ammonium hydroxide, over the escaping vapors. If it 
blackens, the tin itself contains a little sulfide and cannot be used for the most | 
delicate test. Usually, however, the blackening with tin alone is so slight that 
an allowance can be made for it. Now add a little of the substance to be tested 
and a little more tin. A trace of sulfide will cause blackening of the lead 
acetate paper. 


6. Sodium Nitroprusside, Nag{Fe(CN)s(NO)]-2H20, is colored 
reddish-violet by 5 ions, but not by SH ions. Consequently hydrogen 
sulfide itself does not give this reaction, except upon the addition 
of caustic ‘alkali. The reaction is very sensitive, but not so delicate 
as the one with an alkaline solution of a lead salt. 

7. Methylene Blue.—This reaction (which was recommended by 
Emil Fisher *) is the most sensitive of all reactions for detecting the 
presence of hydrogen sulfide. It is particularly suited for detecting 
the presence of traces of. hydrogen sulfide in mineral waters, evep — 
when all other tests give negative results. 





* Ber., 16, 2234. 


HYDROGEN SULFIDE 337 


Treat the solution to be tested for hydrogen sulfide with one tenth of its 
volume of concentrated hydrochloric acid, add a little dimethylparaphenylen- 
diamine sulfate, NH.-CsH,-N(CH;)2-HSO,, from the point of a knife-blade, 
stir it into the liquid, and when it has dissolved, add one or two drops of a 
dilute ferric chloride solution. 

The formation of methylene blue may be expressed by the following equation: 


2[NH, , C.H, ’ N(CHs): , HCl] +6Fet + F4e§= 


Jf Otln —N (CHs)2 
S +6Fet ++NH,++4Ht+Cr. 


—N 
NH; —==N (CHs)2 


Cl 


If only 0.02 mgm. of hydrogen sulfide is present in a liter, the blue color is 
distinctly apparent after half an hour’s standing, while the above tests would 
give negative results. 

If too little hydrochloric acid is present, a red coloration is obtained; this 
is caused by the action of ferric chloride upon a faintly acid solution of di- 
methylphenylendiamine. If considerable hydrochloric acid is present, the red 
coloration does not appear. 


8. Oxidizing Agents, such as the halogens, nitric acid, chromates, 
permanganates, ferric salts, etc., decompose hydrogen sulfide with 
separation of sulfur. 

9. Metallic Silver is blackened by both free hydrogen sulfide and 
soluble sulfides: 


2Ag+H2S+0 (air) =H20+Ag.S; 
2Ag+Na2S+H20+0 (air) =2Na0H+Ag.S. 


If oxygen and water are not present, the above reactions will not 
take place. A piece of bright silver suspended for fourteen hours in 
a sulfur spring showed no sign of darkening until it had been exposed 
~ to the air for a short time. ) 

Absolutely dry hydrogen sulfide, in the presence of absolutely 
dry oxygen, acts upon silver at ordinary temperatures only very slowly; 
it acts instantly if a trace of water is present. 


To detect the presence of sulfur in insoluble sulfides, fuse with a little 
caustic soda (on the cover of a porcelain crucible) to form soluble sodium sulfide: 


NiS+2Na0H =H.,0+Ni0O+Na8. 


Some sulfate is always formed by this treatment; but the aqueous solution 
of the melt will always contain enough alkali sulfide for any of the above tests. 


4 


338 - REACTIONS CF THE ACID CONSTITUENTS 


Behavior of the Sulfides on Ignition 


Most sulfides remain unchanged when heated out of contact with 
the air; arsenic and mercuric sulfides sublime. ’ 

The polysulfides lose sulfur, which sublimes. The sulfides of gold 
and platinum lose sulfur, leaving the metal behind. All sulfides when 
heated in the air give off sulfur dioxide, which can be recognized by 
its odor. 


The Detection of Sulfur in Non-Electrolytes is usually effected by 
heating the substance in a glass tube with metallic sodium (ef. p. 291), and 
testing the aqueous extract of the melt with sodium nitroprusside; or, the resi- 
due in the tube may be treated with dilute hydrochloric acid and the escaping 
gas tested with lead acetate paper for hydrogen sulfide. 


The following method of testing for sulfur is very certain. It 
depends upon the conversion of any sulfur present into the sulfate 
ion which is tested for barium chloride in hydrochloric acid solution. 
(Cf. p. 404). The best way of converting the sulfur into sulfuric 
acid is to heat with concentrated nitric acid in a sealed tube. | (Carius 
method, see Vol. II) or, in the case of difficultly volatile substances 
low in sulfur, by fusion with sodium peroxide in a nickel crucible. 
Since, however, this last oxidation often takes place with explosive 
violence, it is best not to use pure sodium peroxide, but to mix it with 
sodium-potassium carbonate. Mix the substance (from 0.1 to 5 
gms. according to the sulfur content) with ten times as much sodium- 
potassium carbonate and three times as much sodium peroxide and 
heat in a nickel crucible, with the crucible inserted in a disk of 
asbestos board to keep the flame of the gas away from the contents; 
if this precaution is not taken a little sulfate is obtained from the 
sulfur in the gas. After cooling the melt, dissolve it in water, filter, 
acidify with hydrochloric acid and test the filtrate for sulfuric acid 
with barium chloride, | . 


SULFUR, S 


M. Pt. =113.5°—119.5°; B. Pt. 444.5°. 


Occurrence.—Sulfur is found native in voleanic regions in the form of ortho- 
rhombic pyramids, and in the neighborhood of sulfur waters, being formed from 
the oxidation of some of the hydrogen sulfide which escapes. 

Preparation and Properties.—Like the halogens, sulfur is formed by the 
oxidation of its hydrogen compound: 


By heating polysulfides or the sulfides of shes noble metals (gold and plati- 
num), sulfur is also obtained. 

Sulfur exists in three allotropic modifications: 

1. As Orthorhombic Sulfur, with a melting-point of 113.5° C., obtained by 
crystallization from solutions below 95°. 

2. As Monoclinic Sulfur, with a melting-point of 119.5° C., obtained by the 
solidification of molten sulfur. 

3. As Amorphous Sulfur, obtained by quickly cooling the molten sulfur after 
it has become viscous by heating to 250° C., or after it has become a thin a 
after heating to a higher temperature. 

Monoclinic sulfur changes gradually into orthohombic octahedrons; or, 
in other words, the unsymmetrical form changes into the symmetrical form. 
This is a general phenomenon: 

If a substance exists in two or more crystallographic forms, the more symmetrical 
form is always the more stable, and the less symmetrical always has a tendency to 
go over into the more symmetrical form. Thus the unsymmetrical, orthorhombic, 
yellow mercuric iodide is changed, by rubbing with a glass rod, into symmetrical, 
tetragonal, red mercuric iodide (cf. p. 196); and, similarly, the orthorhombic 
form of calcium carbonate, aragonite, changes into hexagonal calcite. 

Both of the crystalline modifications of sulfur are soluble in carbon disul- 
fide; and, by evaporating the solution, the sulfur always recrystallizes in the 
form of octahedrons. Amorphous sulfur is insoluble in carbon disulfide. 

Commercial “‘ flowers of sulfur” is a mixture of crystalline and amorphous 
sulfur, and therefore is only partly soluble in carbon disulfide. 

Sulfur burns in the air to sulfur dioxide, and, in the presence of “ contact 
substances,”’ such as platinum, oxide of iron, chromic oxide, etc., it is burned 
to sulfur trioxide also. Consequently the gas from pyrites burners always 
contains a mixture of the two gases. 

Sulfur is insoluble in water, but soluble in hot caustic alkali, forming a, 
thiosulfate and a sulfide: 


6Na0OH +48 = 3H,0 +. Na.S.0; 4+ 2Na.S. 
By further action of sulfur the Na2S is changed into NaS:, NasSs, ete. 


This reaction is entirely analogous to the formation of hypochlorite and 
chloride by the action of chlorine upon cold dilute caustic alkali: 


2NaOH+Cl, = H.0+Na0Cl+NacCl. 
339 


ole in alkali sulfites, forming a t 


. f ‘ . 


seta -—NaSOstS=NaS.0., 





ACETIC ACID , HC2H302 


Acetic acid is found in the sap of many plants, partly free and 
partly in the form of its potassium or calcium salt. 

It is formed by the dry distillation of wood or by the oxidation 
of alcohol. | 

Anhydrous acetic acid (glacial acetic acid) solidifies below +16° 
C., forming colorless, glistening plates. It has a penetrating odor, 
similar to that of sulfur dioxide, and is miscible with water, alcohol, 
and ether in every proportion. It boils at 118° C. 

The aqueous solution reacts acid. It is monobasic acid, and 
its salts, the acetates, are as a rule readily soluble in water; the silver 
salt is difficultly soluble. 

The most important commercial salts of this acid are sodium 
acetate and lead acetate (sugar of lead), Pb(C2H302)2+3H20. 

Neutral lead acetate dissolves lead oxide with the formation of 
soluble basic salts: 


—_OH 

Pb(C2H302)2+PbO+H20=2Pb_ G57, 3 
—O—Pb—CsH30> 

Pb(C2Hs02)2+2PbO=Pb_ 4: bh 6.7.05" 


The solutions of the soluble basic lead acetates, as. well as that of neutral 
lead acetate, yield precipitates of lead carbonate when acted upon by carbon 
dioxide. For this reason turbid solutions are often obtained when these salts 
are dissolved in distilled water, because the latter frequently contains carbonic 
acid. If a drop of acetic acid is added to the turbid solution the precipitate 
disappears at once. 


REACTIONS IN THE WET WAY 


Use a solution of sodium acetate for the following reactions: 

1. Dilute Sulfuric Acid sets acetic acid free from its salts; it is 
quite volatile and can be recognized by its odor. 

2. Concentrated Sulfuric Acid also sets acetic acid free. If alcohol 
is added at the same time and the mixture warmed, ethyl acetate 


is formed, ) 
HC2H302+C2Hs50H = H20+Ce2Hs-CoH302, 


Ethylacetate 


which can be recognized by its pleasant, fruity odor. 
341 ° 


342 REACTIONS OF THE ACID CONSTITUENTS 


3. Silver Nitrate produces, in fairly concentrated solutions, a — 
white crystalline precipitate of silver acetate (100 parts of water 
dissolve 1.04 parts at 20° C. and 2.52 parts at 80° C.). 

4, Ferric Chloride colors neutral solutions dark brown, and by 
boiling the diluted solution, all the iron is precipitated as basic acetate 
(ef. p. 150). 


REACTIONS IN THE DRY WAY 


All acetates are decomposed on ignition, leaving behind either 
the carbonate, oxide, or the metal itself, and with the evolution of 
combustible vapors and gases. 

The acetates of the alkalies are decomposed into carbonate and 
acetone: 

2NaC2H302 = NagCO3+(CHs3)2CO. 


The acetates of the alkaline earths always leave the metal in the 
form of its oxide, while the acetates of the noble metals leave a resi- 
due of the metal itself. 

Cacodyl Reaction.—If a dry acetate (best an alkali acetate) is 
heated with arsenic trioxide, a very repulsive-smelling and extremely 
poisonous gas is formed, called cacodyl oxide: 


4NaCoH302 + AseOz et 2Na2CO3+[(CH3)2As]20+2CO02 T ° 


In spite of the sensitiveness of this test, it cannot always be 
relied on, for many other organic acids, such as butyric and valeri- 
anic acids, give similar reactions. 


Cyanic AcID, HCNO 


This very unstable acid is obtained by heating its polymer, cyan- 
uric acid, (HCNQ)3; it is a colorless liquid with a very penetrating, 
disagreeable odor, which immediately decomposes in aqueous solution, 
forming ammonium bicarbonate: 

HCNO+2H20 — HCO2NH2+H20 — NH4HCOs. 
Carbamic acid - 

The salts of cyanic acid, the cyanates, are much more stable than 
the free acid, and may be obtained by the oxidation of cyanides. 

By simply fusing potassium cyanide in the air, perceptible quan- 
tities of potassium cyanate are formed. If, however, potassium 
cyanide is heated with oxidizing substances, or those which can be 
readily reduced, it is easy to change the cyanide completely over 
to cyanate. ‘The cyanates of the alkalies are stable in the dry state, 
but take on moisture from the air and are gradually changed into 
alkali bicarbonate and ammonia: 


KCNO+2H20 =KHCO3+NHs ¢. 


Solubility of Cyanates—The cyanates of the alkalies and alkaline 
earths are soluble in water. Silver, mercurous, lead, and copper 
cyanates are insoluble in water. All cyanates are soluble in nitric 
acid. 

REACTIONS IN THE WET WAY 


Use a freshly prepared, cold solution of potassium cyanate for 
these reactions. 

1. Dilute Sulfuric Acid immediately sets cyanic acid free, which 
decomposes into ammonium salt and carbon dioxide: 


CNO-+H*t — HCNO; HCNO+2H20 — NH4*+HCOs;"; 
HCO3-+H* — H20+C02 fT. 


Consequently a strong evolution of carbon dioxide takes place on 
adding the sulfuric acid. The carbon dioxide always contains small 
amounts of undecomposed cyanic acid, which is recognizable by its 
very penetrating odor. The solution will contain ammonium sulfate; 
if it is warmed with caustic soda, ammonia will be given off. 

343 


344 REACTIONS OF THE ACID CONSTITUENTS 


2. Concentrated Sulfuric Acid reacts similarly. 
3. Silver Nitrate precipitates white, curdy silver cyanate, 


KCNO+AgNO3=KNO3+AgCNO, 


soluble in ammonia and in nitric acid (difference from silver cyanide). 

4. Barium Chloride produces no precipitation. 

5. Cobalt Acetate is colored azure blue by a solution of potassium 
cyanate. The blue potassium cobaltocyanate, Ke[Co(CNO)a4], dis- 
covered by Blomstrand,* ‘is formed by this reaction, and is obtained 
in the form of tetragonal crystals of a dark azure-blue color. : 

This blue compound dissolves in water with a blue color. If, however, 
it is subjected to the action of considerable water, the color disappears, the 
double salt being dissociated into its components: 


K,Co(CNO), = Co(CNO).+2KCNO. 


If more potassium cyanate is added to the solution, which has become 
colorless, the blue color reappears. The same result is reached by adding 
alcohol. 

Almost all commercial potassium cyanide contains some cyanate. 

In order to detect the presence of cyanate in the commercial salt, the cyan~ 
ide must first be expelled, for the cobalt test does not take place in the presence 
of cyanide. | 

According to E. A. Schneider,} the test is made as follows: 

Dissolve 3 to 5 gms. of the potassium cyanide in 30 to 50 ec. of cold water 
and pass carbon dioxide into the solution for 60 to 90 minutes; the hydrocyanic 
acid is expelled, and potassium bicarbonate carbonate is formed, while the 
potassium cyanate is not affected perceptibly: 


KCN+H:CO;=HCN f +KHCO;. 


Take 1 ce. of the solution, add 25 cc. of absolute alcohol (to precipitate 
potassium bicarbonate) and filter. Treat the alcoholic filtrate with a few 
drops of acetic acid and then add a few drops of alcoholic cobalt acetate solution. 

If the original cyanide contained 0.5 per cent of potassium cyanate, the blue 
color can be distinctly recognized, but alkali thiocyanates give the same reaction. 





* Journal fir praktische Chemie [2], 3, 206. 
T Ber., 1895, p. 1540. 


HYPOPHOSPHOROUS ACID, H3POz2 


Hypophosphorous acid is obtained by the decomposition of its 
barium salt with sulfuric acid, or of its calcium salt with oxalic acid. 
The acid is monobasic, only one of the hydrogen atoms being replace- 
able by metals. The salts of hypophosphorous acid are obtained 
by boiling phosphorus with dilute alkali, whereby phosphine is given 
off : 

2P4+3Ba(OH)2+6H20 =3Ba(H2PO2)2+2PHs3 f ; 


P1+3KOH+3H20 =3K(H2PO2)+PH3 T ; 


The phosphine thus obtained is spontaneously combustible owing 
to the presence of small quantities of liquid phosphine, P2H4; it is, 
however, mixed with considerable hydrogen because the alkali reacts 
upon the alkali hypophosphite with evolution of hydrogen. See No. 
6, below. 

- Solubility of Pr spenkoucied: —All hypophosphites are soluble in 
water. 

REACTIONS IN THE WET WAY 


1. Dilute , Sulfuric Acid.—No reaction. 

2. Concentrated Sulfuric Acid reacts with hivnaplicenhtea only 
on warming, and is reduced to sulfur dioxide, which can be recog- 
nized by its odor. 

3. Silver Nitrate is reduced to metallic silver, sometimes with and 
sometimes without the evolution of hydrogen, according to the relative 
amounts of the substances reacting: 


2H2PO2 +2Ag*+4H20 — 2H3P04+2Ag+3Hpe fT , 
HePO2 +4Agt+2H20 — H3P04+4Ag+3H". © 


4. Barium Chloride causes no precipitation. 
5. Copper,* Mercury, and Gold Salts are reduced to metal. 
6. Concentrated Caustic Potash.—By boiling with concentrated 





* With copper the reduction may go so far that copper hydride is formed. Cf. 
" Wirz, Compt. rend., 18, 102. 
345 


346 REACTIONS OF THE ACID CONSTITUENTS 
caustic alkali, the hypophosphites are oxidized, with evolution of 
hydrogen, to phosphates: 
H2PO2°+20H™ — POe+2He J. 
7. Nascent Hydrogen (zinc and dilute sulfuric acid) reduces hypo- 
phosphorous acid to phosphine (see Phosphorous Acid). 
REACTIONS IN THE DRY WAY 
By ignition phosphate and phosphine are obtained: 
2H3PO02 — H3P04+PHs f ; 
4NaH2PO2 — NasP207+H20 f +2PHs 7 ; 
2Ca(H2PO2)2 — CazP207+H20 7 +2PH3 7. 


GROUP III 


Silver Nitrate produces a white precipitate, soluble in nitrie acid. 
Barium Chloride does the same. 


SULFUROUS ACID, H2SO; 


— Occurrence and Preparation—Sulfur dioxide, the anhydride of 
sulfurous acid, is found in the exhalations of active volcanoes, and 
is formed by the combustion of sulfur or sulfides in the air. 


S+02=S0O2 fT, 
4FeSo+1102 = 2Fe203+8S02, 


or by the reduction of sulfuric acid on heating with sulfur, sulfides, 
carbon, organic substances, and metals: 


2H2804+8 =2H20+3S02 T ; 
2H2804+C =2H20+CO2 fT +2802 i gr 
2H2804+ Cu =2H20+CuS04+SO02 1 OF 


Mercury, silver, tin, etc., act the same as copper. 
Sulfur dioxide is also formed by the decomposition of sulfites and 
thiosulfates with stronger acids: 


NazSOs-+H2804=Na2S04+H20+S80z f ; 
NaeSe03+ HeS04 = NaeSO4+8+He0+S802 tT ; 


Sulfurous acid may be prepared for laboratory purposes by placing 
a concentrated solution of sodium bisulfite in a flask and allowing 
-concentrated sulfuric acid to drop upon it. A steady stream of sulfur 
dioxide will be evolved without warming. . : 

Sulfur dioxide is a colorless gas, having the penetrating odor 
peculiar to burning sulfur, and is readily soluble in water and alcohol: 
one vol. water at 15° C. dissolves 45.36 vol. SQ2; one vol. alcohol at 
15° C. dissolves 116 vol. SOz. 

The aqueous solution contains sulfurous acid, H2SO3. The acid 
cannot be isolated, as it decomposes on evaporation into water and 
sulfur dioxide; consequently the free acid is known only in aqueous 

347 


348 REACTIONS OF THE ACID CONSTITUENTS 


solution. By neutralization of this solution with alkali hydroxides 
or carbonates, the comparatively stable sulfites are obtained. Jn 
solution, sulfites are gradually oxidized to sulfates by dissolved oxygen. 

Solubility of Sulfites——The sulfites of the alkalies are readily 
soluble in water; the remaining sulfites are difficultly soluble or in. 
soluble in water, but are all soluble in hydrochloric acid. 


REACTIONS IN THE WET WAY 


1. Dilute Sulfuric Acid evolves sulfur dioxide, in the cold, from 
all sulfides, the gas being easily recognized by its odor. 

2. Concentrated Sulfuric Acid reacts in the same way, but much 
more energetically. 

3. Silver Nitrate produces, in neutral solutions of sulfites or in 
an aqueous solution of sulfurous acid, a white crystalline precipitate 
of silver sulfite, 


SO3-+2Agt — AgoSOs, 
soluble in an excess of alkali sulfite, forming sodium silver sulfite: 
AgeSO3-+ Na2SO3 = 2Nal[AgSOs]. 
By boiling this solution the silver is precipitated as a gray metal: 
2[AgSO3l — SOs" +802 7 +2Ag. 


If water containing silver sulfite in suspension is boiled, half the 
silver is reduced to metal, while the other half goes into solution as 
sulfate: 

2Ag2SO3 — =AgeSO4+S8O2 T +Age. 


Silver sulfite is soluble in ammonia and in nitric acid. 

4, Barium Chloride produces no precipitation in an aqueous solu- 
tion of sulfurous acid, but in neutral sulfite solutions white barium > 
sulfite is precipitated, 


SO3-+Ba** — BaSOs, 


readily soluble in cold, dilute nitric acid. By boiling the solution, 
barium sulfate is formed slowly and separates out. As sulfites in 
aqueous solution gradually change to sulfates, commercial sulfites 
usually contain sulfate. In this case the precipitate produced by 
barium chloride in neutral solution contains barium sulfate, which 
is insoluble in dilute nitric or hydrochloric acid. If the barium sul- 
fate is filtered off and the filtrate treated with chlorine or bromine 


SULFUROUS ACID 349 


water, a white precipitate of barium sulfate is formed, provided a sul- 
fite was originally present: 


803-+Cle+Ba*t*+H.0 — 2H+ +2Cl-+BaSO,. 


5. Strontium and Calcium Salts behave similarly to the barium 
salt. 

The sulfites of the alkaline earths vary in their solubilities in 
sulfurous acid and in water. 

Calcium sulfite readily dissolves in an excess of sulfurous acid, 
forming calcium bisulfite: 


CaSO; + H2SO3 = Ca(HSOs) 2. 


On boiling this solution, sulfur dioxide escapes, and calcium sul- 
fite is reprecipitated. | 

The strontium salt also dissolves in sulfurous acid, but more dif- 
ficultly; the barium salt is practically insoluble in sulfurous acid. 


SOLUBILITY OF THE ALKALINE-EARTH SULFITES IN WATER 


One part calcium sulfite dissolves in 800 parts water at 18° C. 
One part strontium sulfite dissolves in 30,000 parts water at 18° C. 
One part barium sulfite dissolves in 46,000 parts water at 18° C. 


Advantage is taken of the difficult solubility of strontium sulfite 
in detecting sulfurous acid in the presence of thiosulfuric acid (which 
see). . 
6. Lead Salts precipitate white lead sulfite, soluble in cold dilute 
nitric acid; but on boiling the solution lead sulfate is precipitated. 

7. Sodium Nitroprusside and Zinc Sulfate——If a neutral sulfite 
solution is treated with a dilute solution of sodium nitro-prusside, a 
faint pink coloration is produced. If, however, considerable zine sulfate 
is added, the coloration becomes a distinct red. The reaction is still 
more sensitive if a little potassium ferrocyanide is added, a red pre- 
cipitate being formed (difference from thiosulfuric acid). This reac- 
tion, although very delicate, is not so reliable as the precipitation 
with strontium chloride. 

Sulfurous acid is a strong reducing agent. 

8. Iodine Solutions are decolorized by sulfurous acid: 


SO3-+H.0+1. > 2H*+21I-+S80.4". 


9. Acid Potassium Permanganate Solutions are also decolorized, 
sulfuric and dithionic acids being formed in varying amounts accord- 
ing to the temperature and concentration. 


350 REACTIONS OF THE ACID CONSTITUENTS | 


Under certain conditions the reaction can take place according 
to the following equations: 


2MnO4~+6S8037-+8Ht — 2Mn*t+4804>+820¢67+4H20. 


Under other conditions, however, the sulfurous acid can be com- 
pletely oxidized to sulfuric acid. Consequently, sulfurous acid cannot 
be determined quantitatively by means of potassium permanganate. 

10. Chromic Acid is reduced to green chromic salt: 


2CrO4~ + 3503 + 10H* =< 2Crt + Fe 3804->+ 5H20. 


11. Mercuric Chloride is unaffected by sulfurous acid at ordinary 
temperatures; but, on boiling, it is reduced to mercurous chloride, 


2H¢gCle +S8037°+H20 — 2Ht+ 2Cl- +8047 +Hge2Cle : 


and on adding more sulfurous acid, the mercurous salt is reduced 
to gray metal. 

12. Mercurous Nitrate is immediately acted upon by free sul- 
furous acid and by alkali sulfite solutions, with the formation of a 
black precipitate. 

13. Gold Solutions are also reduced. 

14. Nascent Hydrogen reduces sulfurous acid to hydrogen sul- 
fide, which can be recognized by its odor and by its turning lead 
acetate paper black. The reduction is best effected with zinc and dilute 
hydrochloric acid. 


REACTIONS IN THE DRY WAY 


The sulfites of the alkalies, when heated out of contact with the 
air, are changed to sulfate and sulfide: 


4Na2SO3 ais 4 3Na2S0.,+Na.S. 


By heating an alkali sulfite in the closed tube this reaction takes 
place, and there is no sublimate of sulfur (difference from thiosul- 
fates). If the melt is treated with hydrochloric acid after cooling, 
hydrogen sulfide is given off freely. 

The remaining sulfites are changed, on being heated out of con- 
tact with the air, into sulfur dioxide and oxide or metal: 


CaSO3=CaO+S0Osz J ; 
2Ag2SO3 = 4Ag+2802 T +Oe T ; 


If a sulfite is heated with sodium carbonate on charcoal, sodium 
sulfide is formed. If the melt is placed upon a bright silver coin and 


SULFUROUS ACID 351 


.moistened with water, the silver is blackened, owing to the formation 
of black silver sulfide (Hepar reaction) : 


2Na2S03+3C =3CO2 T +2Na2S, 
and 
NaeS+2Ag+H20+0=2Na0H+Ag2S. 


This Hepar reaction takes place with all sulfur compounds, and 
therefore shows simply the presence of sulfur. The oxygen required 
in the above reaction is obtained from the atmosphere. 


CARBONIC ACID, H2COz 


Like sulfurous acid, pure carbonic acid does not exist; it is known 


only in aqueous solution. Its anhydride, COs, is formed by the com- 
bustion of carbon and of carbonaceous matter of all kinds, and is found 
therefore very widely distributed in nature (in small amounts in the 
atmosphere, and in enormous amounts in volcanic regions, streaming 
out from fissures in the earth). 


Pure air contains 0.35-0.40 per cent of CO:. In dwelling places the amount 
increases considerably, owing to breathing and other forms of combustion. 
If 3 or 4 per cent is present, as is the case in mines sometimes, breathing 
becomes difficult, and the miners’ lamp begins to burn faintly; while when 8 
to 10 per cent is present the lamp goes out. 


Carbon dioxide occurs also in many mineral waters, and (in the 


liquid state) is found enclosed in quartz, feldspar, etc. As carbonate 
it exists in enormous quantities as limestone, marble, aragonite, dolo- 
mite, etc. Carbon dioxide is a colorless, odorless, slightly acid-tast- 
ing gas, with sp.gr. 1.52. Being, therefore, one and one-half times as 
heavy as air, it can be poured from one vessel into another. Carbon 
dioxide does not support combustion; a burning candle goes out in 
air containing 8 to 10 per cent of this gas.* 


Carbon dioxide gas is not very soluble in water. At 15° and 760 mm. pres- 
sure, 1 liter of water dissolves its own volume of gas. At higher temperatures 
it is less soluble and it is easy to expel carbon dioxide from a solution. In the 
aqueous solution, the following equilibria exist 


CO.+H,0 < H,CO;  H++HCO;"; 
HCO, Ht+Coz". 


The ionization constant for the primary ionization of carbonic acid is 0.0.3 and 
for the secondary ionization it is 0.0:>7. Aécording to these values, the primary 
ionization of carbonic acid takes place to about 0.1 per cent and there is present 
in a liter of saturated carbon dioxide solution at 15° only about 0.0006 equiv- 
alent of hydrogen ions. The secondary ionization upon which the quantity 
of CO;~ ions present depends, takes place only to a negligible extent. 





* Carbon dioxide not only fails to support combustion, but it tends to prevent 
it; hence its use in fire extinguishers. Being formed by the combustion of car- 
bon, the mass-action principle shows that it will tend to stop the reaction much 
better than an inert gas. 


352 


CARBONIC ACID. 353 


When'a strong acid is added to the solution, even the primary ionization of 
carbonic acid is repressed almost completely. Similarly, when hydrogen ions 
are added to a carbonate, carbon dioxide is formed, even with acetic acid, and 
the carbon dioxide is easily expelled by heating. 


The salts of carbonic acid, the carbonates, are formed: 

1. By passing carbon dioxide gas into a solution of a metallic 
hydroxide: 
: 2Na0OH +COe2 = HeO +NaeCOz ; 


Ba(OH)2+CO2=H20+BaCOsz. 


2. By the action of carbon dioxide upon cyanides, sulfides, and 
borates of the alkalies and alkaline earths. 

3. By the ignition of salts of organic acids (cf. p. 80). 

An illustration of the preparation of large amounts of carbonate 
is the production of potash by burning parts of plants (wood, for 
example, or the residue from the manufacture of beet sugar, the latter 
being particularly rich in potassium salts). 

Solubility of Carbonates.—Of the normal carbonates, only those 
of the alkalies are soluble in water; and their aqueous solutions react 
alkaline, owing to hydrolytic dinoaipodiiiow: 


Na2CO3+H20 @ 2Nat+OH-+HCO3-. 


The aqueous solution of the carbonates of the alkalies, therefore, 
behaves as if it were a solution of caustic alkali and alkali bicar- 
bonate. 

Many carbonates dissolve in an excess of carbonic acid, forming 
bicarbonates, particularly the alkaline-earth carbonates: 


CaCO3+ H2CO3 = Ca(HCOs)e. 


By boiling a solution of calcium bicarbonate, the latter is decom- 
posed into water and carbon dioxide, and calcium carbonate is repre- 
_ cipitated: 

Ca(HCQs)2 =H.20+C0O2 T +CaCOsz. 


Nearly all samples of drinking-water contain calcium or mag- 
nesium bicarbonate; they become turbid, therefore, on boiling (boiler 
scale). Dilute, cold mineral acids decompose all carbonates with 
effervescence (due to evolution of carbon dioxide gas). 

The native carbonates of magnesium and iron (magnesite, siderite, 
and dolomite) do not effervesce if a lump of the mineral is treated 
with cold dilute mineral acids, but when reduced to a fine powder 
they are more readily acted upon; on warming, all carbonates dissolve 
readily. 


354 REACTIONS OF THE ACID CONSTITUENTS 


REACTIONS IN THE WET WAY 


1, Dilute Sulfuric Acid decomposes all carbonates with efferves- 
cence; except with magnesite, siderite, and dolomite, the reaction 
takes place in the cold. 


As the atmosphere always contains carbon dioxide gas, particularly in a 
laboratory where many gas flames are burning, considerable caution is neces- 
sary in testing for a small quantity of carbonic acid. In the first place, if the 
substance does not effervesce with acid, there is no use in making the test. 
Moreover, carbon dioxide is odorless and it is absurd to think that effervescence 
implies the presence of a carbonate when the escaping gas has the odor of hydro- 
gen sulfide or sulfur dioxide. 

A simple method of testing is as follows: Place about 1 em. of the powdered 
solid in a test-tube, cover it with about 10 cc. of water and boil for about a 
minute. This serves to expel the air from the substance and from the water. 
Add a little 6-normal hydrochloric acid and watch closely to see if there is any sign 
of effervescence. If there is effervescence notice whether the escaping gas has 
any odor. . Dip two stirring rods into barium hydroxide solution, place one of the 
rods between the second and third finger and the other between the third and 
fourth and hold one of the rods inside the test-tube, without touching the sides, 
and the other rod outside. If a carbonate is present, the barium hydroxide on 
the rod inside the test-tube will become turbid faster than that on the rod held 
in the air. This test is of no value, however, in the presence of a sulfite which 
will also cause barium hydroxide to become turbid. 

The above test naturally fails to detect traces of carbonate with certainty, 
but it is satisfactory for all ordinary work except when a sulfite is present. 
To detect traces of carbonate, even in the presence of a sulfite, the following 
procedure is useful. Fit a 100 cc. flask with a rubber stopper containing two 
holes. Through one hole insert a small dropping funnel so that it reaches nearly 
to the bottom of the flask. Through the other hole insert a right-angled glass 
tube which serves to lead the escaping gas to a second flask, containing about 
5 gms. of chromic acid anhydride, CrOs;, dissolved in a little water and 25 ce. of 
6-normal sulfuric acid. This will serve to oxidize any sulfurous acid or hydrogen 
sulfide that may be set free. Arrange the tubes in this second flask so that the 
gas passes down to the bottom of the solution and leaves the flask through 
an exit tube which just reaches below the rubber stopper. Connect this flask . 
in the same way with a third flask and connect the exit tube from this flask with 
a drying tube filled with soda-lime, to prevent carbon dioxide getting in from 
the air. Place the powdered substance in the first flask and cover it with 
25 cc. of water. With the third flask empty, conduct a stream of air free from 
carbon dioxide through the apparatus for ten minutes while heating the 
water in the first flask. This is best accomplished by applying suction at 
the end of the train and drawing the air through soda-lime in a drying tube 
which is placed in a rubber stopper that fits the neck of the dropping funnel. 
Instead of soda-lime, caustic potash solution (1 : 2) may be used to:remove the 
carbon dioxide from the air. When the carbon dioxide has been expelled from 
the apparatus, close the stop-cock in the dropping funnel, take away the flame. 
and quickly add about 25 cc. of barium hydroxide solution to the last flask. 
With the apparatus all connected, introduce 25 cc. of 6-normal sulfuric acid 


CARBONIC ACID — 355 


into the dropping funnel and allow it to run slowly into the flask containing the 
substonce. If not enough gas is evolved to produce a turbidity in the barium 
hydroxide solution, apply gentle suction as before, after a slight vacuum has 
been produced, carefully open the stop-cock of the dropping funnel, with the 
soda-lime tube in place, and gradually heat the sulfuric acid to boiling. Con- 
tinue drawing air through the apparatus for fifteen minutes if necessary. In 
case a very slight turbidity is obtained in the barium hydroxide solution it is 
best to run a blank on the apparatus with all the reagents, and then repeat 
the experiment. If the test does not give a decided result, the presence of 
carbonate should never be reported. 


2. Concentrated Sulfuric Acid reacts in the same way as dilute 
sulfuric acid, only more violently. 

3. Silver Nitrate precipitates white silver carbonate, which be- 
comes yellow on the addition of an excess of the reagent. On boiling 
with considerable water, the carbonate is partly decomposed into 
brown silver oxide and carbon dioxide; but the carbonic acid is not 
expelled completely except by heating to 200°. Silver carbonate is 
very soluble in ammonia and in nitrie acid. 

4. Barium Chloride precipitates white, voluminous barium car- 
bonate, in the cold, which gradually on standing, but more quickly 
on warming, becomes crystalline and denser. 


Behavior of Carbonates on Ignition 


The carbonate of the alkalies melt with but slight decomposi- 
tion. Barium carbonate is not decomposed on charcoal before the 
blowpipe, and does not melt; only at a white heat is it decomposed 
into infusible barium oxide and carbon dioxide. All remaining 
carbonates are decomposed at the temperature of the blowpipe into 
oxide and carbon dioxide. The oxides of the noble metals are decom- 
posed further into metal and oxygen. 


PERCARBONIC ACID, HeC20¢ 


Free percarbonic acid is not known, but its potassium salt is stable in the 
dry state. When exposed to moisture it decomposes into hydrogen peroxide 
and potassium bicarbonate: | 


K.C.0,+2H,0 = H,0O.+ 2KHCO:3. 


If the salt, which is characterized by its pale blue color, is placed in con- 
siderable cold, dilute sulfuric acid, it dissolves with evolution of carbon dioxide 
and formation of potassium acid sulfate and hydrogen peroxide. The solution 
gives all the characteristic reactions of hydrogen peroxide. 


To distinguish between percarbonic acid and hydrogen peroxide, dissolve 


10 gms. of potassium iodide in water and add 0.1 to 0.3 gm. of the finely 
powdered substance. If potassium percarbonate is present, iodine is at once 
liberated: 

C.0.~+2I- —> 2CO;~+I2._ 


If only potassium bicarbonate and hydrogen peroxide are present, the liberation 
of iodine will take place much more slowly. 


Silver nitrate and barium chloride when treated with a percarbonate, give — 


white precipitates which are soluble in dilute nitric acid. 


eee 


~~) oe A 


BORIC (BORACIC) ACID, H3BO3 


Occurrence.—Boric acid is found native as sassolite; in the form 
of its sodium salt, as borax or tinkal, NazB407-10H20; as boracite, 
2Mg3Bs015+MgCle; and in many silicates, such as axinite, tour- 
maline, datolite, etc. 

_ Crystallized borie acid forms colorless plates, with a mother-of- 
pearl luster, which are soluble in water (100 parts water dissolve 
4 parts of boric acid at 15°, and 33 parts at 100°). The aqueous 
solution reacts acid, and is a poor conductor of electricity. 

By heating boric acid to 100°, it loses one molecule of water and 
is changed to metaboric acid, HBOz. The latter loses more water 
when heated to 160°, forming pyroboric acid, H2Bs407; which, on 
ignition, loses all its water, being changed to the anhydride of boric 
acid, boron trioxide, which remains as a difficultly-volatile, hygro- 
scopic glass. 

The salts of boric acid, the borates, are derived from the meta- 
and pyroboric acids. The salts of the ortho acid, H3BOs3, are not 
known in the pure state. 

In a few exceptional cases boron acts as a metal, forming B(HSO;);, (BO).SO, 


BF;,BPOx,, etc. The last compound is insoluble in water and dilute acids, but 
dissolves readily in caustic alkalies. 


Solubility of Borates—The borates of the alkalies dissolve in water, 
and the solution reacts alkaline. 

A concentrated solution of borax behaves as if it contained sodium 
metaborate, free boric acid, and a small amount of caustic alkali: 


Na2eB407+3H20 @ 2NaBO2+2H3BO3; 
NaBOo+ 2HOH @ NaOH+H3BOs. 


The more dilute the solution, the greater the extent to which the 
hydrolysis represented by the second equation will take place; so 
that a very dilute solution of borax will react as if it contained simply 
sodium hydroxide and free boric acid. 

A solution of an alkali borate will behave differently towards 
reagents, therefore, according to its concentration and temperature. 
The remaining borates are difficultly soluble in water, but readily 
soluble in acids and in ammonium chloride solution. 

357 


358 REACTIONS OF THE ACID CONSTITUENTS 


REACTIONS IN THE WET WAY 


For these reactions use a borax solution. 

1. Dilute Sulfuric Acid.—No reaction. 

2. Concentrated Sulfuric Acid.—No visible reaction. Most bo- 
rates are decomposed by sulfuric acid, setting free boric acid, and 
the latter is capable of coloring the non-luminous gas-flame with 
a characteristic green tinge. 

If, therefore, a little solid borate is placed in the loop of a platinum wire, 
moistened with concentrated sulfuric acid, and heated at the edge of the 
Bunsen flame, the characteristic green coloration will be noticed. 

A great many natural silicates containing boric acid, when tested in the 
above manner, will not give this flame coloration. To produce this coloration 
moisten the mineral with hydrofluoric acid, place a little of it in the loop of a 


platinum wire, and heat at the outer edge of the flame; if a borate is present the 
latter will be colored distinctly green, owing to the formation of volatile boron 


fluoride. 
The presence of copper or of barium interferes with this test. 


3. Concentrated Sulfuric Acid and Alcohol.—If an alkali or alka- 
line-earth borate is treated in a platinum crucible with methyl alcohol, 
then with concentrated sulfuric acid, the mixture stirred and the 
alcohol lighted, a green-bordered flame will appear, due to the forma- 
tion of boric acid methyl ester, B(OCHs)3. 

4, Silver Nitrate produces, in moderately-concentrated, cold borax 
solutions, a white precipitate of silver metaborate: 


NazgB407 + 3HeO + 2AgNOsz = 2NaNO3 ~1. 2H3BO3 +2AgBOo. 
On warming, a brown precipitate of silver oxide is obtained: 
2AgBO> +3H.O = 2H3BO3 +Ag20. 


From very dilute solutions, in the cold, silver nitrate produces a 
brown precipitate oi silver oxide. 

Silver borate is soluble in ammonia and in nitric acid. 

5. Barium Chloride produces, in fairly concentrated solutions, a 
white precipitate of barium metaborate: 


NaeBs407 +BaClo +3H:2O a 2NaCl +- 2H3BO3 +Ba(BOzg)2, 


soluble in an excess of barium chloride and in ammonium chloride. _ 

6. Calcium and Lead Salts behave similarly to barium chloride. 

7. Turmeric.—If a piece of turmeric paper is placed in a solution 
of free boric acid, apparently no change will take place unless consider- 
able quantities of boric and sulfuric acids are present, but if the 
paper is dried, it becomes reddish brown. If the brown paper is 


BORIC ACID 359 


again dipped in the solution of boric acid, the color remains; which 
is also true if the paper is dipped in a dilute sulfuric or hydrochione 
acid solution (difference from the alkali test with turmeric paper) 
If the reddish-brown paper is moistened with caustic soda or potash 
solution, the paper becomes bluish black; or, if only a small amount 
of boric acid is present, grayish blue. 


The shade and intensity of the color varies both with the amount of tur- 
meric and with the amount of boric acid; with a very little boric acid, turmeric, 
and very dilute caustic soda solution a nearly pure violet color is obtained and 
with considerable boric acid and stronger alkali, a greenish-black color. 

Borate solutions, when acidified with dilute hydrochloric acid, give the above 
reaction. This sensitive and convenient test for boric acid must be used with 
caution, for acid solutions of zirconic, titanic, tantalic, niobic, and molybdie 
acids also color turmeric paper brown. 

The reaction is much more sensitive if, instead of using the turmeric paper 
itself, an alcoholic solution of turmeric is used. Place 2 or 3 drops of the yellow 
solution in a porcelain dish, add the solution to be tested for boric acid, acidify 
with acetic acid, and evaporate.to dryness on the water-bath. If-as much as 
0.02 mgm. of BO; is present, the residue is colored a distinct reddish brown, 
while 0.002 mgm. suffices to cause a visible reaction. (IF. Henz.) 


8. Mercuric Chloride produces a red precipitate of basic mercuric 
salt. If considerable free boric acid is present there is no precip- 
itation. 


Behavior of Borates on Ignition . 


The hydrated borates of the alkalies melt with effervescence, 


forming a colorless glass. 

This glass has the property of dissolving many metallic oxides 
when heated, whereby the often very characteristically colored meta- 
borates are formed (borax beads); thus copper oxide is dissolved, 


forming a blue glass: 
NazBsO7+ Cu0 = 2NaBO2+Cu(BO2)s. 


If this bead is heated in the reducing flame (i.e., with carbon) 


two things can happen: 
(a) The colored cupric salt is reduced to colorless cuprous salt: 


4NaBO2+2Cu(BO2)2 +C= CO+Na2B407+2NaB02+Cuz(BO2)2. 


(b) The cupric salt is reduced to. metallic copper, so that the bead 
appears reddish brown and opaque: 


4NaBOo+2Cu(BO2)2+C =CO2 T +2Na2Bs074+ 2Cu. 


OXALIC ACID, H2C204 


Occurrence and Preparation.—Oxalic acid occurs, in the form of 
its acid potassium and calcium salts, in the sap of many plants. 

It is prepared in large amounts by fusing sawdust with caustic 
alkali. The resulting potassium salt is precipitated with milk of 
lime, forming the insoluble calcium salt; and the latter is decom- 
posed with sulfuric acid. Oxalic acid is also formed by the oxida- 
tion of innumerable organic substances (such as sugar, starch, cellu- 
lose (paper), by means of concentrated nitric acid. 

It crystallizes from aqueous solutions in the form of colorless mono- 
clinic prisms, H2C204-2H20. ; 

By allowing the hydrated acid to stand over sulfuric acid the 
water is lost, and the anhydrous acid remains, which, when heated — 
to about 150° C., sublimes, forming needles. If heated still higher 
it is completely decomposed into water, carbon dioxide, and carbon 
monoxide: 


HeC204= H20+COez FT +CO 7. 


The crystallized, hydrated acid is soluble in water, alcohol, or 
ether: 100 parts water at 20° dissolve 11.1 parts oxalic acid; 100 
parts alcohol at 15° dissolve 33.2 parts oxalic acid; 100 parts ether 
at 15° dissolve 1.5 parts oxalic acid. 

Oxalic acid is a fairly strong, dibasic acid, and forms neutral 
and acid salts, e.g., potassium oxalate, K2C204; potassium bin- 
oxalate, KHC204; potassium tetroxalate, KHC204, HeC204:2H20. 

Solubility —The oxalates are mostly insoluble in water, with the 
exception of the oxalates of the alkalies and of magnesium. In an- 
excess of an alkali oxalate many of the insoluble oxalates dissolve. 
Thus ink spots and rust spots can often be removed from clothing 
by means of a solution of oxalic acid or of potassium tetroxalate: 


Fe203+6H2C204 > 2Ha[Fe(C204)]s +3H20. 


Oxalice acid is also a good solvent for ferric phosphate. All oxalates 
dissolve readily in mineral acids. 


REACTIONS IN THE WET WAY 


A solution of ammonium oxalate may be used for the following 
reactions: 


360 


OXALIC ACID 361 


1. Dilute Sulfuric Acid.—No reaction. 

2. Concentrated Sulfuric Acid, on warming, acts as a dehydrating 
agent, causing the evolution of equal volumes of carbon monoxide 
and carbon dioxide; the latter burns with a blue flame: 


H2C204 > H20+CO T +COe f.. 


In the presence of manganese dioxide, all oxalates evolve CO, with dilute 
H.SO,: ; 
H,C,0,+MnO,+H,SO, =MnSO,+2H,0+2CO, fT. 


In the same way CO, is given off by the action of KMnQ, and dilute H,“O, 
at about 60° C.: 


2KMn0,+5H.C,0,+3H,80, =K,80,+2MnSO,+8H,0+10C0, ft. 
3. Silver Nitrate precipitates white, curdy silver oxalate, 
C204 +2Ag” =AgeC20z, 


almost insoluble in water, but readily soluble in ammonia and in 


- nitric acid. 


4. Barium Chloride precipitates white barium oxalate soluble in 
oxalic and acetic acids. 

5. Calcium Chloride precipitates white calcium oxalate, insoluble 
in oxalic acid, ammonium oxalate, and acetic acid, but readily soluble 
in hydrochloric and nitric acids. It is the most insoluble of all oxalates. 

6. Lead Salts precipitate white lead oxalate, soluble in nitric acid. 


Behavior of Oxalates on Ignition 


~All oxalates are decomposed on ignition with slight carboniza- 
tion. The oxalates of the alkalies and alkaline earths are changed 
to carbonates, with evolution of carbon monoxide. Stronger ignition 
causes the formation of more or less oxide, the alkaline earth carbonate 
being quantitatively changed to oxide by heating over the blast 
lamp. The oxalates of the noble metals, and of iron, nickel, cobalt, 
copper, etc., leave the metal itself; the oxide is formed in the case of 
other metals. 


TARTARIC ACID, H2C4H.10, 


Occurrence.—Tartaric acid occurs partly free and partly as its acid 
potassium salt in many fruit saps, particularly in that of the grape. 

The free acid crystallizes in clear, monoclinic prisms, without 
water of crystallization. Its aqueous solution is optically active, 
turning the plane of polarized light to the right. 


Three other modifications of this acid exist, possessing the same chemical 
formula, but differing in their physical properties. One of these turns the plane 
of polarized light to the left, and the other two are optically inactive. 


Tartaric acid is very readily soluble in water (100 parts water 
dissolve 132 parts of tartaric acid at 15°) and alcohol, but it is insol- 
uble in ether. The salts are called tartrates. 

Solubstlity—The neutral alkali tartrates are very soluble in water, 
as also is acid sodium tartrate, while the acid potassium and the acid 
ammonium tartrates are difficultly soluble in water. 

The remaining tartrates are difficultly soluble in water, but all 
dissolve, more or less readily, in neutral alkali tartrate solution, form- 
ing complex salts. 

The most important commercial salts of this acid are ‘‘ cream 
of tartar,’ KHC4H40¢, “ Rochelle salt,’ KNaC4H40¢, and “ tartar 
emetic,’ K(SbO)C4H40¢. 


REACTIONS IN THE WET WAY 


A solution of Rochelle salt (sodium potassium tartrate) may be 
used for these reactions. 

1. Dilute Sulfuric Acid.—No reaction. 

2. Concentrated Sulfuric Acid causes carbonization on warming, 
with evolution of sulfur dioxide. 

3. Silver Nitrate produces no precipitation in a solution of free 
tartaric acid, but in the solution of a neutral tartrate, a white, curdy 
precipitate is immediately formed, 


C4H40¢6 + 2Agt — AgeC4H10.,, 


readily soluble in nitric acid and in ammonia. By warming the 
ammoniacal silver solution, metallic silver is deposited. This very 
362 


TARTARIC ACID } 363 


important reaction for the detection of tartaric acid is performed 
in the following manner: 


Treat the pure tartrate solution with silver nitrate solution until no further 
precipitation takes place, then add dilute ammonia drop by drop until the pre- 
cipitate just dissolves. Place the test- tube containing the solution in water 
which has been heated to 60-70° C. After standing for about fifteen minutes, 
the silver will be deposited in the form of a beautiful mirror on the sides of the 
test-tube. This very delicate reaction cannot be performed with certainty in 
the presence of other acids. In this case the tartaric acid should first be 
' precipitated as potassium acid tartrate. Concentrate the solution to a small 
volume, add a little solid potassium carbonate, acidify with strong acetic 
acid and stir the cold solution vigorously; a precipitate of potassium acid 
tartrate will form at once if considerable tartrate is present. Filter off the 
precipitate, wash with a little cold water, and dissolve it in as little caustic 
soda solution as possible. In this way a solution is obtained which will 
readily give the silver mirror on the addition of silver nitrate and treatment as 
above. 

If no precipitate is formed -on the addition of the acetic acid, add a little 
alcohol, which causes the precipitate to form more readily. Filter off the 
precipitate, wash with diluted alcohol, dry, dissolve in dilute sodium hydrox- 
ide, and treat as above. If the alcohol is not removed by drying, a mirror is 
sometimes formed when tartaric acid is absent. 


4. Calcium and Barium Chloride.—If to a concentrated solution 
of neutral alkali tartrate, in the absence of ammonium salts, calcium 
chloride solution is added drop by drop, a white amorphous pre- 
cipitate is formed which redissolves, forming soluble calcium tartrate 
anions: 


2C4H4106 +Ca — [Ca(C4H40¢)o]-. 


Only after the addition of enough calcium chloride to decompose 
completely the alkali tartrate is a permanent precipitate formed, 
which at first is flocculent, but soon becomes crystalline, consisting 
of neutral calcium tartrate: 


[(Ca(C4H40¢)2]> +Cat* — 2CaCsH4O¢. 


In dilute solutions the first addition of calcium chloride often produces no 
precipitation; but after standing some time (or more quickly on rubbing the 
sides of the test-tube with a glass rod) the crystalline precipitate is deposited, 
CaC.H,O.+4H.0. Calcium tartrate is very difficultly soluble in water; 100 
parts water at 15° C. dissolve 0.0159 part of the crystalline salt, and 100 parts 
boiling water dissolve 0.0285 part of the salt. The precipitate is soluble in 
acetic acid (difference from calcium oxalate) and also in a solution of concen- 
trated caustic alkali (free from carbonate), probably forming a complex ion: 


2CaC,H,0O,+20H- — H,0+[(CaCyH.Q¢),0]-. 


364 REACTIONS OF THE ACID CONSTITUENTS 


On boiling this solution, calcium tartrate is reprecipitated in the form of a 
voluminous gelatinous precipitate, which again goes into solution on cooling.. 
The presence of ammonium chloride retards the formation of the calcium tar- 
trate, but does not prevent it; after standing some time, the precipitate settles 
out in the form of a heavy crystalline powder (difference from citric acid). 


5. Potassium Salts produce no precipitation in neutral solutions 
of alkali tartrates; but if the concentrated solution is acidified with 
acetic acid, a precipitate of crystalline potassium acid tartrate is 
formed upon stirring vigorously: 


C4H40¢6-+Kt+Ht — KHC4H10¢. 


Potassium acid tartrate is difficultly soluble in water (100 parts water 
dissolve 0.45 part of salt) and in acetic acid, but is readily soluble in mineral 
acids or in caustic alkali and alkali carbonate solutions. The precipitate may be 
dissolved in a little ammonia and the above test with calcium chloride obtained. 

If a concentrated solution of free tartaric acid is treated with potassium 
chloride, a precipitate of potassium acid tartrate is formed in spite of the pres- 
ence of the hydrochloric acid which is set free. From dilute solutions the pre- 
cipitate appears only after adding sodium acetate (cf. p. 79). : 

The presence of considerable boric acid greatly interferes with the formation 
of the potassium acid tartrate. In such cases add ammonium chloride to the 
concentrated solution in a test-tube, then some calcium chloride solution, and 
rub the sides of the glass. with a stirring rod. If only a little tartaric acid is 
present it may be necessary to let the solution stand twenty-four hours. Filter 
off any precipitate that may form, wash it two or three times with alcohol, and - 
dissolve the calcium tartrate in a little freshly-prepared potassium hydroxide 
solution (1:5). Filter and heat the filtrate to boiling. Amorphous calcium 
tartrate should be precipitated. Another way to identify the tartrate in the 
first precipitate of calcium tartrate is as follows: Place the washed precipitate 
in a test-tube, add a crystal of silver nitrate, a few drops of 6-normal ammo- 
nium hydroxide, and place the test-tube in hot water; a distinct silver mirror 
should form if a tartrate is present. 


6. Lead Acetate produces in neutral solutions a white, flocculent 
precipitate of lead tartrate, easily soluble in nitric acid and in ammonia. 

7. Magnesia Mixture. If a concentrated tartaric acid solution is 
treated with an excess of magnesia mixture, 10 ce. of strong ammonia 
and a volume of alcohol equal to that of the solution, then, after 
shaking and allowing to stand twelve hours, the tartaric acid is pre- 
cipitated quantitatively as crystalline, basic magnesium tartrate 
insoluble in 50 per cent alcohol (difference from malic and succinic 
acids) : 


C4H106—>+2Mg* *+20H~+H20 — Mgo(OH)2 (C4H40¢)- H20. 


Filter off the precipitate obtained in the above test, wash it with 50 per 
cent alcohol, dry and transfer the precipitate to a test-tube with the aid of a 


' TARTARIC ACID 365 


glass rod. Add a little silver nitrate solution, a slight excess of 6-normal 
ammonium hydroxide and heat to about 60°. A mirror should form if a tar- 
trate is present. 


REACTIONS IN THE DRY WAY 


If tartaric acid is heated to 135° C., it melts, and on stronger 
ignition it is decomposed, leaving a residue of carbon and giving off 
empyreumatic odors (smell of burnt sugar). 

The alkali tartrates are also decomposed by ignition, leaving a 
residue of carbon and alkali carbonate, which effervesces on treat- 
ment with acid. | 

Ammonium tartrate leaves a residue of carbon, which does not 
effervesce on treatment with acids. The tartrates of the alkaline 
earths leave behind a mixture of carbon and carbonate; on very 
strong ignition the latter is changed to oxide. 

The tartrates of those metals whose oxides are reduced by car- 
bon are left in the form of metal (Ag, Pb, Fe, Ni, Co, ete.). 


CITRIC ACID, H3C 6Hs5O07 


Citric acid is found in nature in the juices of many fruits. It is 
a tribasic acid, readily soluble in water and in alcohol, but difficultly 
soluble in ether. Its salts are called citrates. 

Solubility —The citrates of the alkalies are soluble in water, and 
form, with the insoluble citrates of the heavy metals, very soluble 
complex salts, whose solutions are not precipitated by alkali hydrox- 
ides, alkali «carbonates, ammonia, etc. 


REACTIONS IN THE WET WAY 


A solution of potassium citrate may be used. 

1. Dilute Sulfuric Acid.—No reaction. 

2. Concentrated Sulfuric Acid on being heated with a citrate, as 
with most a-hydroxyacids, causes formic acid, HCO2H, to be formed, 
which then’ breaks down into water and CO; at the same time some 
acetone dicarbonic acid, (CH2)2CO-(COe2H)s2, is formed and this breaks 
down into acetone, (CH3)2CO, and COs. 


H3C.Hs07 =? HCO2H+(CH2)2CO(CO2H )2, 
HCO2H — H20+C0 T, (CH2)2CO’COsH)2 —.(CH3)2CO+2C02 T . 


A part of the citric acid is carbonized and this causes reduction of the - 
sulfuric acid so that some, SQOz2 is evolved. 

3. Silver Nitrate produces in neutral solutions a flocculent pre- 
cipitate of silver citrate, AgsCe6HsO7, readily soluble in nitric acid 
and in ammonia. On heating the ammoniacal solution to 60° C., 
no silver mirror is formed; but on heating the solution to boiling, 
the silver is gradually deposited. 

4. Barium and Calcium Chloride give no precipitation in neutral 
solutions (difference from tartaric acid). If, however, caustic soda 
solution is added to the solution which contains an excess of calcium 
chloride, a flocculent precipitate of tertiary calcium citrate is at once 
formed, insoluble in caustic alkali, but readily soluble in ammonium 
chloride. On boiling the solution in ammonium chloride, crystalline 
calcium citrate is precipitated, which is now insoluble in ammonium 
chloride. 

366 


CITRIC ACID 367 


5. Lime Water in excess produces no precipitation in cold solutions 
of neutral citrates; on boiling, there is formed a flocculent precipitate 
of calcium citrate, which almost entirely redissolves on cooling. 

_ 6. Lead Acetate precipitates from solutions of the free acid, and 
those containing neutral salts, amorphous Pb3(CsH;07)2-H2O. 


7. L. Stahre’s Test for Citric Acid.* To the solution of free citric acid in 
water, or to the solution of a citrate in very dilute sulfuric or nitric acid (not 
hydrochloric) add 2 to 5 drops of tenth-normal permanganate solution and heat 
a short time at 30° to 40° (the solution must not boil!). As soon as the solution 
is colored brown, or becomes turbid by the precipitation of a little manganese 
dioxide, add 1 or 2 drops of ammonium oxalate solution and about 1 ce. of 
10 per cent sulfuric acid, which will clear up the solution. Now, add a few 
drops of bromine water, and a distinct, crystalline precipitate of pentabromace- 
tone will be obtained. The bromine water may also be added before the per- 
manganate solution and sometimes the results are better. 

This test is so sensitive 0.38 mgm. of citric acid in 1 ec. of water can be 
detected. 

The experiment succeeds in the presence of tartaric, malic, oxalic, sulfuric 
and phosphoric acids, except that a little more permanganate is required. 

In the Stahre test the following reactions take place: 

(a) The permanganate oxidizes the citric acid to acetonedicarboxylic acid 
with evolution of carbon dioxide: 


€H;CsH;0;+2Mn0, +€Ht — 5CO, T +2Mn*t++£H.0+5(CH,)2,CO(CO.H).. 


(6) The acetonedicarboxylic acid reacts with bromine, forming penta- 
bromacetone: 


(CH,),CO(CO.H).+5Br, a 2CO, T +5HBr+C.HBr,CoO. 


If the permanganate is allowed to act longer upon the citric acid the acetone. 
dicarboxylic acid is converted gradually into acetone, the reaction taking 
place more quickly on boiling: 


(CH,).CO(CO.H)2 > 2CO, T +(CH:;).CO. 


Acetone itself is not brominated as readily as the acetonedicarboxylie acid, 
and for this reason care should be taken not to let the temperature rise above 
40° during the treatment of the citric acid with permanganate. 

The citrates on treatment with bromine, without previous oxidation with 
permanganate, will also give pentabromacetone: 


K;C.H;0; + 6Brz =3KBr + 3CO, T + 4HBr+ C.HBr;CO, 


8, Mercuric Sulfate.—Denigés’ reagent.t Dissolve 5 gms. HgO in 100 ee. 
of water and 20 ee. cone. H.SO,. Treat the solution of the citrate with 1/20 
as much reagent and heat to boiling, then add a few drops of 0.1N KMnO, 
solution. A white crystalline precipitate is formed. 





*L. Sraure, Z. anal. Chem., 36 (1897), 195; also ALFRED WOuLK, ibid., 41, 
94 (1902). 
{ Comptes rend., 188, 32; Z. anal. Chem., 88, 718 (1899); and 40, 121 (1901). 


368 REACTIONS OF THE ACID CONSTITUENTS 

The precipitate has the composition: Hg;0.80,-2[(CH2),CO(CO:).JHg and 
is a mixture of basic mercuric sulfate and the mercuric salt of acetone dicar- 
boxylic acid. The reaction is very sensitive and enables one to detect 0.5 gm. 
of citric acid dissolved in a liter of water. The reaction is, however, not pecu- 
liar to citric acid but is shown by many other ketonic compounds, | 


REACTIONS IN THE DRY WAY 


The citrates, on ignition, behave exactly like the tartrates 


PHOSPHOROUS ACID, H3PO3 


-Formation—By the slow combustion of phosphorus in the air 
phosphorus trioxide is formed which, as it is the anhydride of phos- 
phorous acid, reacts with cold water to form the acid: 


P203+3H20 =2H3PO3. 


Phosphorous acid is formed much more readily by the action of 
water on the trihalogen compounds of phosphorus: 


PCl3+3HOH =3HCl-+ H3P0Os. 


The hydrochloric acid is removed by evaporation, and the last 
traces of uncombined water by heating to 180°. If the mass is then 
allowed to cool, it solidifies to a crystalline, hygroscopic substance 
which melts at 70°. 

By neutralizing the solution of phosphorous acid with bases; the 
phosphites are obtained. It is never possible, however, to replace 
more than two of the hydrogen atoms with metal; so that phosphorous 
acid is considered a dibasic acid. Certain organic compounds are 
known, however,-which are derived from tribasic phosphorous acid, 
H3POs3. 

Solubility Only the phosphites of the alkalies are soluble in 
water, but they are all soluble in acid. 


REACTIONS IN THE WET WAY 


A solution of sodium phosphite should be used. 

1. Dilute Sulfuric Acid.—No reaction. 

2. Concentrated Sulfuric Acid causes no reaction in the cold; on 
heating, the phosphorous acid reduces the sulfuric acid to sulfurous 
acid, easily recognized by the odor of burning sulfur, 


H3PO3-+H2S04 = H3P01+H20+80z2 f . 


3. Silver Nitrate produces at first a white precipitate of silver 
phosphite, 


HPO3-+2Ag* — AgoHPOs, 
369 


370 REACTIONS OF THE ACID CONSTITUENTS 


which in the case of a concentrated solution is changed in the cold 
to metallic silver; while in dilute solutions this reduction takes place 
only on warming: 


AgoHPO3; -}- H2O = H3P04 +2Ag. 


4. Barium Chloride precipitates white barium phosphite, soluble 
in all acids. 

5. Lead Acetate precipitates white lead phosphite, insoluble in 
acetic acid. 

6. Mercuric Chloride is te reduced phosphorus acid in 
the cold, but more quickly on warming, to mercurous chloride: 


2H¢gClo+H3P03+H20 = H3P01+2HCl+Hg2Ch. 


If the phosphorous acid is present in excess, the reduction in the 
hot solution (not in the cold) goes further, and gray metallic mercury 
is deposited: 


HgeCle a. H3PO3 +H.O = HsPO4 -f- 2HCl +2H¢g. 


7. Nascent Hydrogen (zinc and sulfuric acid) reduces phosphor- 
ous acid to phosphine: 


H3P03+6H =3H20+PHs} . 


If the phosphine is allowed to act upon a concentrated solution 
of silver nitrate (1:1), or better still, upon the solid silver nitrate, 
the latter is colored yellow, as with arsine: 


PH3+6AgNO3 = PAgs3-3AgNO3+3HNOs. 


By the addition of water this yellow compound is decomposed 
with separation of grayish-white silver: 


PAgs . 3AgNOz3 +3H2O = 3HNOs: +H3P03 +6Ag. 


The phosphorous acid is, however, immediately oxidized by the 
nitric acid to phosphoric acid: 


3H3P03+2HNO3 = He0+2N0+3H3POs4. 
The mixture of phosphine and hydrogen burns with an emerald- 


green flame. 


' 8. Sulfurous Acid is reduced by phosphorous acid to hydrogen 
sulfide: 


3H3PO03+ H2S03 =3H3P04+ HS f . 


PHOSPHOROUS ACID 371 
9. Concentrated Potassium Hydroxide Solution changes a phos- 
phite to phosphate, with evolution of hydrogen, 
KeHPO3+ KOH = K3P04+ He fF , 
but with dilute caustic potash the hydrogen evolution is very slight. 


REACTIONS IN THE DRY WAY 


By ignition, phosphorous acid (like hypochlorous acid) is changed 
at the cost of its own oxygen to the higher compound, while the oxidiz- 
ing part of the acid is reduced to its hydrogen compound: 


3HCI1O = HClO3+2HCl; 
4H3PO3=3H3P04+PHs f . 
The phosphates behave similarly: 
SNazHPO3 =4NazgP04+ NasP207+H20+2PHs 7 . 


METAPHOSPHORIC ACID, HPO; 


The monobasic metaphosphoric acid is obtained by treating phos- 
phorous pentoxide with cold water, 


P205;+H20 =2HPOs, 
and also by the strong ignition of orthophosphoric acid: 
H3P04= H20+HPOs. 


Metaphosphoric acid is a colorless, glassy, hygroscopic mass. 
On boiling its aqueous solution, or slowly in the cold, it adds water to- 
the molecule, and is changed to orthophosphoric acid: 


HPO3+ H20 = H3P04. 


The metaphosphates are readily obtained by heating the mono- 
metallic salts of orthophosphoriec acid, 


NaH2PO4=Hz20 fT +NaPOs, 
or by igniting sodium ammonium phosphate: | 
NaNH4HPO4=H20 , +NH3 T +NaPOs. 


The meta salts are changed into orthophosphates by boiling the 
aqueous solution in the presence of mineral acid. 

Solubility.—The metaphosphates of the alkalies and of magnesium 
are soluble in water; the remaining salts are difficultly soluble or 
insoluble in water, readily soluble in nitric acid, and in an excess of 
metaphosphoric acid or an excess of alkali metaphosphate. 


REACTIONS IN THE WET WAY 


Sodium metaphosphate is used for the following tests: 

1. Sulfuric Acid causes no visible reaction. . 

2. Silver Nitrate precipitates white silver metaphosphate, soluble 
in ammonia and in nitric acid: 


PO3--+Agt — AgPOs. 


3. Barium Chloride precipitates voluminous barium metaphos- 
phate, soluble in an excess of sodium metaphosphate, from which 
372 


METAPHOSPHORIC ACID 373 


solution ammonia causes no precipitation. Barium sodium dimeta- 
phosphate (or a similar polymetaphosphate) is probably formed. 

4. Magnesium Salts cause no precipitate from moderately dilute 
solutions, even on boiling (difference from orthophosphoric acid). 

5. Ammonium Molybdate produces no precipitate in the cold; 
but, on boiling the acid solution, metaphosphoric acid is changed to 
orthophosphorie acid and the characteristic precipitate of ammonium 
phosphomolybdate is formed. . 

6. Albumin Solution is coagulated by an aqueous solution of the 
free acid (difference from pyro- and orthophosphoric acids), but not 
by a solution of alkali metaphosphate, except on the addition of acetic 
acid. 

7. Nascent Hydrogen does not reduce metaphosphoric acid (differ- 
ence from phosphorous acid). | 


Behavior on Ignition 


The alkali metaphosphates, on being fused, form a glassy mass, 
which has the property of dissolving many metallic oxides, forming 
orthophosphates with characteristic colors. (See Phosphoric Acid.) 
By fusion with soda, orthophosphates are formed from metaphos- 
phates. 


PYROPHOSPHORIC ACID, H4P207 


The tetrabasic pyrophosphoric acid is formed by heating ortho- 
phosphoric acid to 213°. It is a soft, glassy mass, readily soluble in 
water; and in solution it gradually adds water to the molecule. and is 
changed to phosphoric acid, the change taking place quickly on 
boiling the solution. 

The salts of pyrophosphoric acid, the pyrophosphates, are obtained 
by igniting the dimetallic phosphates: 


2Na2HP O4 = He2O +NaaP207. 


Solubility—The pyrophosphates of the alkalies are soluble in 
water; the remaining pyrophosphates are difficultly soluble or insol- 
uble in water, but are all soluble in acids, and some are soluble in an 
excess of sodium pyrophosphate. 


REACTIONS IN THE WET WAY 


A solution of sodium pyrophosphate is used for these tests. 

1. Sulfuric Acid.—No reaction. 

2. Silver Nitrate gives a white, curdy precipitate, soluble in am- 
monia and in nitric acid. 

3. Barium Chloride causes a white, amorphous precipitate, coer 
in acids. 

4. Magnesium Chloride produces a white precipitate which is 
soluble in an excess of the magnesium salt, as well as in an excess 
of sodium pyrophosphate. By boiling this solution a_ precipitate 
is formed, which does not disappear on cooling. 

5. Ammonium.Molybdate produces no precipitation in the cold; 
but, on warming, yellow ammonium phosphomolybdate is precipi- 
tated. 

6. Albumin is not coagulated by free pyrophosphoric acid (differ- 
ence from metaphosphoric acid). ; 


BEHAVIOR IN THE DRY WAY 


All pyrophosphates on being fused with sodium carbonate are 
changed to orthophosphates: 


NasP207+Na2CO3=CO2 T +2NagPO04. 
374 


IopIc ACID, HIO; 


Occurrence.—In sea-water and in Chili saltpetre as potassium 
iodate. 

Formation.—By oxidizing iodine with fuming nitric acid or by 
the action of chlorine upon iodine suspended in water: 


312+10HNO3=6HIO3+10NO T +2H20; 
Iz+6H20+5Cle = 10HC1+2HI103. 


The most important iodate, KIOs3, is obtained by the action of 
iodine upon a slightly acid solution of potassium chlorate: 


5KC10O3+3I12+3H20 = 5KIO3+ HIO3+5HCI. 


Iodates are also formed by the action of iodine upon alkali hy- 
droxide solutions: 


3le+6KOH = 5KI+ KIO3+3H20. 


In alkaline solutions iodides are oxidized to iodates by hypo- 
chlorites and potassium permanganate. 

Solubility —The -iodates of the alkalies are soluble in water, but 
the remaining iodates are difficultly soluble or insoluble. 


REACTIONS IN THE WET WAY 


1. Sulfuric Acid.—Neither dilute nor concentrated sulfuric acid 
decomposes iodic acid; but if reducing substances are present at the 
same time (such as hydriodic acid, hydrogen sulfide, ferrous salts, : 
etc.), the iodic acid is reduced, with separation of iodine: 


I03~+5I-+6H* — 3H20+43le. 


2. Silver Nitrate precipitates white, curdy silver iodate, AgIOs, 
readily soluble in ammonia, but difficultly soluble in nitric acid. 

3. Barium Chloride precipitates white barium iodate, difficultly 
soluble in hot water (100 parts of boiling water dissolve 0.6 part of 
the salt), and only slowly soluble in dilute hydrochloric or nitric acids. 

4. Lead Acetate precipitates lead iodate, difficultly soluble in 
water and only slightly soluble in nitric acid. 

375 


- 


376 REACTIONS OF THE ACID CONSTITUENTS | 


5. Reducing -Agents. 
(a) Hydriodic acid reduces iodic acid, with separation of iodine: 


H1I03+5HI=3H20+3le. 


If the solution is concentrated, the iodine separates out as a brown 
powder; dilute solutions are colored yellow. The iodine may be 
absorbed, with a reddish-violet color, by shaking the aqueous solution 
with chloroform or carbon disulfide. 

(b) Sulfurous acid also causes separation of iodine, unless a large — 
excess of sulfurous acid is added: . 


2103~+5S037+2H*t > 5804 +H20+]e, 
IO3_+3803° — 3804-4 T-. 


The reduction takes place according to the last reaction, when 
three molecules of sulfurous acid are present to one of iodate; sul- 
furous acid reacts with free iodine (cf. p. 349). 

(c) Zinc dust (or, better, Devarda’s alloy) reduces neutral iodate 
solutions to iodide. : 


REACTIONS IN THE DRY WAY 


Heated on charcoal the iodates deflagrate, but not so strongly 
as the chlorates; they are all decomposed on being heated, some 
with and some without the separation of iodine. Thus all neutral 
iodates of the alkalies are easily decomposed into iodide and oxygen, 
while the biiodates set free iodine at the same time: 


2KIO3 =2KI+302T ; 
4/KIO3-HIO3] =4KI+1102 fT +2H20+42Io. 


GROUP IV 


Silver Nitrate produces in neutral solutions a colored precipitate, 
soluble in nitric acid. 

Barium Chloride also produces a precipitate which is soluble in 
nitric acid. 


PHOSPHORIC ACID, H3P04 


Orthophosphoric acid is obtained by the oxidation of phosphorus 
by means of nitric acid, or by boiling the meta- and pyro-phosphoric 
acids with water. It is a tribasic acid, and forms salts in which either 
one, two, or three of its hydrogen atoms are replaced by metals (cf. 
p. 10), e.g., NaH2PO4, NazHPO«4 and NagPOs.. 

Solubility—The phosphates of the alkalies are soluble in water, 
and so are the primary salts of the alkaline earths. The secondary 
phosphates of the alkaline earths are very difficultly soluble, while 
the corresponding tertiary phosphates (as well as all other phos- 
' phates) are insoluble. All phosphates dissolve in acids (cf. p. 47). 


REACTIONS IN THE WET WAY 


Use a solution of disodium phosphate for these reactions. 

1. Sulfuric Acid, dilute or concentrated, produces no visible change. 

2. Silver Nitrate produces a yellow precipitate of silver phosphate 
(difference from meta- and pyrophosphoric acids), 


2HPOs" +3Ag*=H2POs. +AgsPOu, 


readily soluble in nitric acid and in ammonia. The precipitate, there- 
fore, can be formed only in neutral solution. 
3. Barium Chloride precipitates white, amorphous barium phos- 
phate: 
HPO.z-+Bat * — BaHPOs.. 


_ In the presence of ammonia the less soluble tertiary salt is precipitated: 
2HPO4-+3Bat +1LONH3 oo 2QNHat+ Baz (PO4)e. 


The barium phosphates (as well as those of the other alkaline 
earths) are easily dissolved by acids, even acetic acid (difference from 
377 


378 REACTIONS OF THE ACID CONSTITUENTS 


aluminium and ferric phosphates). From these acid solutions, am- 
monia reprecipitates the phosphate. 


In pure water H.PO,~ is dissociated to about 0.1 per cent, but in the presence 
of acetic acid to a much less extent. BaHPO, in contact with water furnishes 
more HPO; ions than does H.PO, in the presence of acetic acid, therefore 
BaHPQ, dissolves. Similarly Bas(PO,)2 dissolves in order to establish equilib- 
rium between Ht and PO; ions. Aluminium phosphate is less soluble than 
barium phosphate and therefore requires a stronger acid to dissolve it. 

By adding ammonia the hydrogen ions are neutralized and PO; ions are 
obtained in equilibrium with NH,* ions, and Ba;(PO,)2 is reprecipitated. 


4. Magnesia Mixture (an aqueous solution of ammonium chloride, 
ammonia, and magnesium chloride) precipitates from very dilute 


solutions white, crystalline magnesium ammonium phosphate, 
MgNH4P04+6H20, 


HPO. +Mg*tt+NH3 — MgNHyPO,, 


which is soluble in all acids, but practically insoluble in dilute 23 
per cent ammonia. This is a very sensitive reaction (cf. p. 75). 

5. Ferric Chloride.—If a solution of sodium phosphate is treated 
with ferric chloride, a yellowish-white precipitate of ferric phosphate 
is formed: | 

HPO,-+Fet ++ = Ht+FePQ,. 


Hydrogen ions are formed in this reaction and, as might be expected from 
the mass-action principle, the precipitation of the phosphoric acid is not quan- 
titative unless the greater part of the hydrogen ions are removed. This may 
be accomplished by adding ammonium acetate, as the hydrogen ions must 
form non-ionized acetic acid to be in equilibrium with the acetate ions (cf. 

. 46). 
2 HPO,-+(C.H;0.-+Fet tt — HC.H;0.+FeP O,. 


Moreover, if the reaction takes place in a boiling, dilute solution, the excess 
of the iron can be precipitated as basic ferric acetate. If the solution is filtered 
hot, a filtrate is obtained which is free from iron and from phosphoric acid. If, 
however, it cools, some of the iron goes back into solution and, as ferric phos- 
phate is appreciably soluble in ferric acetate solution, some of the phosphoric 
acid also goes into solution. 

Since ferric phosphate, unlike the phosphates of the alkaline earths, is 
insoluble in acetic acid, it is evident that phosphoric acid may be removed 
from a solution of alkaline earth phosphate in acetic acid by adding ferric 
chloride, an excess of soluble acetate and boiling. 

To accomplish this, dissolve the phosphate in as little hydrochloric acid as 
possible, add ammonium carbonate until a slight permanent precipitate is 
obtained and dissolve the precipitate by adding one or two drops of 6-normal 
hydrochloric acid. Add an excess of ammonium acetate and ferric chloride, 
drop by drop, until the solution above the yellowish-white precipitate of ferric 
phosphate is colored distinctly brown by colloidal ferric hydroxide. Dilute 


PHOSPHORIC ACID 379 


with considerable water, heat to boiling and filter while hot. To detect phos- 
phoric acid in the precipitate dissolve it in nitric acid, evaporate the solution 
to a small volume and treat with ammonium molybdate solution; a yellow, 
crystalline precipitate of ammonium phosphomolybdate proves the presence 
of phosphoric acid. Or, dissolve the iron precipitate in hydrochloric acid, add 
2 gms.of tartaric acid to prevent the precipitation of iron, add ammonia in excess 
and then some magnesium-ammonium chloride mixture. A white precipitate 
of magnesium ammonium phosphate shows the presence of phosphoric acid. 


6. Ammonium Molybdate, in large excess, precipitates from nitric 
acid solutions in the cold on standing (more quickly on slightly warm- 
ing) a yellow, crystalline precipitate of ammonium phosphomolybdate: 


H3P04 + 12(NHa4) 2M004+21HNO3= 
= (NH,)3PO4-12M003+21NHiNO3+12H20. 


This reaction is analogous to the reaction with arsenic acid (cf. page 231), 
except that the arsenic compound is formed quickly only at the boiling tem- 
perature. The presence of ammonium nitrate greatly facilitates the formation 
of this precipitate. 

Ammonium phosphomolybdate is readily soluble in alkalies and in ammonia, 


(NH,);PO,-12M00,+280H- — 3NH.*+HPO+12Mo00,-+1 1.0, 


also in an excess of alkali phosphate solutions, forming conipoulids which 
contain less molybdenum. It is, therefore, always necessary to prevent the 
formation of such compounds by the addition of a large excess of ammonium 
molybdate. 

Detection of Phosphorus in Iron and Steel.—Phosphorus is present in iron 
and steel as iron phosphide, but only to a slight extent (usually less than 0.1 
per cent). To detect the phosphorus it is necessary to oxidize it to phosphoric 
acid and then use one of the above reactions. As, however, very small amounts 
of phosphorus are present, it is necessary to start with a large amount of the 
original substance in order to obtain a perceptible phosphorus test. It is best to 
proceed as follows: Dissolve 5 to 10 gms. of the iron or steel in 60 cc. of 6-normal 
nitric acid,* evaporate the solution to dryness and ignite over a free flame (with 
constant stirring) until no more red fumes are given off. All organic matter is 
thereby destroyed, silicic acid is dehydrated and the oxidation of the phosphorus 
to phosphoric acid is completed. After cooling, dissolve the oxides in 50 ec. of 
12-normal hydrochloric acid (warming gently), evaporate off the excess of acid, 
dilute and filter off the silica. In the filtrate all the iron and all the phosphoric 
acid will be found, and the latter may be detected by either the molybdate or 
the magnesia-mixture reaction. To detect the phosphoric acid according to the 
former method, evaporate to dryness the filtrate obtained after the removal 
of the silica, dissolve the residue in as little 6-normal nitric acid as possible, 
add 50 ec. of ammonium molybdate solution and 15-20 cc. of a 75 per cent 





* If the iron were dissolved in HCl or H2SO,, part or even all of the phosphorus 
would escape as phosphine. Nitric acid oxidizes nearly all of the phosphorus to 
phosphoric acid. 


380 REACTIONS OF THE ACID CONSTITUENTS 


ammonium nitrate solution, heat the mixture gently, shake or stir vigorously 
and allow it to stand an hour. A yellow, crystalline precipitate shows the 
presence of phosphorus. 

To detect the phosphorus according to the magnesia-mixture method, it is 
necessary first to remove the greater part of the iron. Neutralize the hydro-— 
chloric acid filtrate with ammonia, add a saturated solution of sulfur dioxide 
and boil the solution, whereby the previously dark-colored solution is either 
decolorized or becomes a light green. Add 20 cc. of 12-normal hydrochloric 
acid, and boil the solution until the excess of sulfur dioxide is expelled. By 
this operation all the ferric salt is reduced to ferrous salt. Add a few drops of 
chlorine water (which forms a little ferric salt), neutralize with ammonia and 
dilute to about a liter; add 3 cc. of a saturated solution of ammonium acetate, 
5 ec. of acetic acid, and heat the solution to boiling. All the ferric salt and all 
the phosphoric acid will be precipitated in the form of ferric phosphate and 
basic ferric acetate, while the greater part of the iron remains in solution as 
ferrous salt. Filter off the light brown precipitate through a small plaited 
filter, wash it with hot water, and dissolve in dilute hydrochloric acid. Evapo- 
rate the solution almost to dryness, add 2 gms. of citric (or tartaric) acid (which 
should be dissolved in as little water as possible), add an excess of ammonia, and 
precipitate the phosphoric acid by the addition of magnesia mixture. A white, 
crystalline precipitate shows the presence of phosphoric acid. 


7. Lead Acetate precipitates white lead phosphate, nearly insol- 
uble in acetic acid: 


2HPO4 + 3Pb**+ 2C2H302° — 2HC2H302+Pb3(PO4)o. 


8. Nascent Hydrogen does not reduce phosphoric acid (difference 
from phosphorous and hypophosphorous acids). 

9. Metastannic Acid.—If metallic tin is added to a nitric acid 
solution of phosphoric acid, or a phosphate, the tin is changed to 
metastannic acid, which unites with the phosphoric acid, forming 
an insoluble compound (probably a complex phospho-stannic acid). 
This reaction is often used to separate phosphoric acid from other _ 
metals. 

10. Mercurous Nitrate precipitates from solutions which are 
almost neutral, white mercurous phosphate, soluble in nitric acid 
but insoluble in acetic acid. 


REACTIONS IN THE DRY WAY 


The tertiary salts of the alkalies melt without decomposition; 
the secoridary salts lose water and are changed to pyrophosphates 
while, the primary salts form a glassy metaphosphate. 

The so-called “salt of phosphorus,” or “ microcosmic salt,” 
NaNH,HPO.+4H20, which is much used as a reagent, loses water 


PHOSPHORIC ACID 381 


and ammonia on being fused, forming a clear glass of sodium meta- 
phosphate: 


NaNH.HPO,-4H20 =5H20 ¢ +NHs f +NaPOs. 


If the salt is heated in the loop of a platinum wire, a clear bead 
is obtained—the so-called “‘ salt of phosphorus ”’ bead. 

Just as metaphosphoric acid unites with water, on boiling its 
solution, forming orthophosphoric acid, 


HP O3 -}- H2O = H3POsz, 


so sodium metaphosphate dissolves, at the fusion temperature, a 
great many metallic oxides, forming characteristically colored ortho- 
phosphates, 


| NaPO3+Cu0 = NaCuP0Og (blue bead), 
which may be changed in the reducing flame to metaphosphate again: 
NaCuPO4+C=CO JT +Cu+NaP0Os. 


Brownish red opaque bead 





Many anhydrous phosphates are reduced by heating with mag- 
nesium to phosphides, which, on being breathed upon, give the 
peculiar odor of phosphine: 


Cas3(POs)2+8Mg =8Mg0+CasP2; 
CazP2+6H20 =3Ca(OH)2+2PHs f . 


PHOSPHORUS, P. At. Wt. 31.04. Mol. Wt. Py =124.16. M.P.=44.1° 


Occurrence.— Phosphorus is found in nature only in the form of 
phosphates, of which calcium phosphate is the most important. It 
occurs as apatite, Cas(PO4)3(Cl,F), in hexagonal crystals, and in an 
impure state as phosphorite, which is used extensively as a fertilizer. 
Calcium phosphate is also an important constituent of bones and the 
seeds of plants. 

A very interesting occurrence of phosphorus is pyromorphite 
(ef. p. 205), isomorphous with apatite, vanadinite, and mimetesite. 

Properties—Phosphorus exists in four allotropic forms: (a) As 
ordinary or colorless phosphorus. (b) As red, crystalline phosphorus. 
(c) As bright red phosphorus. (d) As black, crystalline phosphorus. 

Ordinary phosphorus is poisonous, is colorless when pure (it be- 
comes yellow on exposure to the light, and is coated with a layer of 
red phosphorus), melts at 44° C., and ignites at 60° C. in the air, 
so that it must be kept covered with water, in which it is insoluble. 
It is readily soluble in carbon disulfide, and slightly soluble in ether. 
It is easily oxidized by nitric acid to phosphoric acid: 


3P4+20HNO3+8H20 = 12H3P04+20NO f. 


The colorless phosphorus, but not the red modification, is oxidized 
to hypophosphorous and phosphorous acids by exposure to moist 
air. This causes the characteristic phosphorous odor, and, in the 
dark, a pale green luminescence. If phosphorus vapors, or phos- 
phine, are allowed to act upon most silver nitrate paper, the latter 
‘is blackened, on account of the formation of silver phosphide and 
metallic silver. The reaction probably takes place in this way: First, 
the phosphorus reacts with water to form phosphine and hypophos- 
phorous acid, 


P4+6H20 =3H3P02+PHs f ,* 
which then react with the silver nitrate: 
H3P02+2H20+4AgNO3 =4HNO3+H3P04+4Ag; 
PH3+3AgNO3 =3HNO3+AgsP. 





* Phosphorus and water by themselves do not react in accordance with this 
reaction, but it seems probable that they do in the presence of silver nitrate. 
382 


PHOSPHORUS 383 


This exceedingly sensitive reaction for colorless phosphorus was 
discovered by Scheuer.* It is a decisive test only when no other 
substance is present, such as HeS, H3As, HgSb, formaldehyde or 
formic acid, which is capable of reacting with silver nitrate. 

Red phosphorus is crystalline (hexagonal, rhombohedral), and 
is formed by heating ordinary phosphorus to about 250° out of con- 
tact with the air. It is not poisonous, is insoluble in carbon disul- 
fide, and does not ignite until heated to 256°. It is non-luminous 
in the dark, does not oxidize in the air, but is readily oxidized by nitric 
acid to phosphoric acid. 

Light-red phosphorus is obtained, according to Schenk,{ by heat- 
ing a solution of white phosphorus in phosphorus tribromide for 
hours with a return-flow condenser. The phosphorus which then 
separates is of a light-red color, is not poisonous, but enters into reac- 
tion so readily that it is easily distinguished from red phosphorus. 
It dissolves in concentrated potassium or sodium hydroxide with a 
stormy evolution of phosphine, the reaction taking place even more 
readily than with white phosphorus. When covered with ammonia, 
it blackens. 

Black phosphorus is obtained when red phosphorus and lead are 
heated together in a sealed tube to a red heat and the mass treated 
with dilute nitric acid after it is cold; the lead dissolves and leaves 
the phosphorus as black phosphorus. By heating to 360° it is changed 
to ordinary phosphorus again. 

Phosphorus is found in a great many organic substances. In 
order to detect its presence, the compound is heated in a sealed tube 
with fuming nitric acid, which destroys the organic matter and oxidizes 
the phosphorus to phosphoric acid (detected by any of the above 
reactions). 

Arsenious, arsenic, and chromic acids, which also belong to this 
group, have already been described on pp, 225, 229 and 135, 


Mitscherlich Test for White, Poisonous Phosphorus { 


This sensitive test is based upon the luminescence of white phosphorus when 
exposed to moist air in the dark. It is used to detect phosphorus in cases of 
poisoning. . 

Procedure.—Place the food residues, or finely-cut pieces of the body, in the 
liter flask K (Fig. 26) and add enough water to form a thin paste. Then, while 
shaking, add tartaric acid to slightly acid reaction, in order to combine with 





* Ann. Chem. Phys., 112 (1859), 214. 
+ Ber., 36, 979 (1903). 
tJ. pr. Chem., 66, 238 (1855). 


384 REACTIONS OF THE ACID CONSTITUENTS 


any ammonia present. Connect the tube R with the flask and heat the contents 
of the latter to boiling. Carry out this operation in a dark room. As the 
vapors condense in the tube A, a greenish luminous zone is visible even when only 
a few milligrams of phosphorus are present. If larger quantities of phosphorus 
are at hand, the distillate in the flask B contains tiny globules of phosphorus 
which, by gently heating and rotating the liquid, can be made to collect into a 
larger drop; the aqueous solution also contains phosphorous acid which can be 
detected by the method of Blondlot-Dusart (see below). 

If, therefore, the luminosity is apparent during the distillation in the dark, 
then the presence of white phosphorus is probable but not certain, because 


| 























Fig. 26. 


phosphorus subsulfide, P.8;, sometimes used as a substitute for phosphorus in 
the manufacture of matches, will often cause luminescence in the Mitscherlich 
apparatus,* especially if a little zine oxide is added to the liquid in K to com- 
bine with H.S, which tends to prevent the luminescence. 

If the luminescence is not apparent, it is not certain that white phosphorus 





*T. Mar and F. Scuarrer, Ber., 1903, 870; L. Vianon, Bull. soc. chem. [3], 33, 
805 (1905), and ScnenKk and Scuarrr, Ber., 1906, 1522. The author wishes to 
state that all commercial preparations of P,S; do not show luminescence in the 
Mitscherlich apparatus. Thus a sample from Kahlbaum did not show the slightest 
luminescence by boiling with water or with concentrated salt solution. It was 
very pure and contained 55,82 per cent P and 44,14 per cent 8, 


PHOSPHORUS 885 


is absent; traces of ammonia, hydrogen sulfide, alcohol, ethereal oils and un- 
saturated hydrocarbons, interfere with the test. In such cases it is advisable 
not to stop distilling too soon, as it often happens that the interfering substances 
will distill over and then the luminescence will appear. In case no: luminescence 
is noticed, examine the distillate. Treat a part of it with strong chlorine water, 
evaporate to small volume on the water-bath and then test for phosphoric 
acid. Cf. pp. 378 and 379. 

According to J. Peset,* the luminosity is very distinct in the flask if the liquid 
is heated to boiling, allowed to cool somewhat, and again boiled. In this way 
0.004 mgm. of phosphorus can be detected. 


Detection of Phosphorus and Phosphorous Acid according to Blondlot- 
Dusart 


This beautiful method is based upon Dusart’s observation that hydrogen 
containing phosphine, when allowed to flow from a tube provided with a plat- 






















TANS -——. 
Li 
enP IOS ‘ ) 
: | 
| 
| 
| 


ouecbaeaeiad 


Seeahy 


aha 





a8 
Lx) A 
Y) 


| TEPPER AHHH 11111 








Fia. 27. 


inum tip, will burn with a flame having an emerald-green core. The green color 
is particularly apparent upon holding a cold porcelain dish in the flame. 

Since phosphorus, phosphorous acid and hypophosphorous acid (not phos- 
phorie acid) are easily reduced to phosphine by zinc and dilute sulfuric acid, 
it is merely necessary to pass the evolved gas through a tube with a platinum 
tip, and light it when the air is all expelled; the merest trace of phosphorus is 
recognized by the green color. 


* Z. anal. Chem., 48, 35 (1909). 
+ L. Dusart, Compt. rend., 48, 1126 (1856), and BLuonptot, J. pharm. chim. 
(3| 40, 25 (1854), 





386 REACTIONS OF THE ACID CONSTITUENTS 


Inasmuch as organic substances can prevent the appearance of the green 
flame, the phosphorus is first separated from it as follows: Place the solution 
containing the phosphorus, or the distillate obtained by the Mitscherlich test, 
in a gas-evolution flask, add zine (free from phosphorus) and dilute sulfuric acid 
(1 : 7), and pass the evolved gas into a neutral solution of silver nitrate; ~ 
if phosphorus is present a black precipitate of silver phosphide is obtained, 
which, if hydrogen sulfide was present, may contain silver sulfide. Filter off 
this precipitate, wash it well with water, and place it in the Blondlot apparatus 
(Fig. 27). In the 500 cc. Woulfe bottle, W, generate hydrogen by means of 
zinc, free from phosphorus, and dilute sulfuric acid (1: 7). After the air is 
entirely expelled from the apparatus, close the pinch-cock a, which causes the 
acid to rise into the reservoir 7' (a bottle with the bottom cut off). Now open 
the cock a wide enough to permit a steady stream of hydrogen to pass out 
from the delivery tube, which is made of potash-glass and is provided with a 
platinum tip.* The flame from the lighted gas should not be too large. 

If the flame shows no green luminescence in the dark when a porcelain dish 
is held in it, then the hydrogen gas is free from phosphorus and can be used — 
for the test. Rinse the black silver precipitate through 7' into the bottle W. 
If the precipitate contained phosphorus, the core of the flame becomes green, 
particularly noticeable upon holding a porcelain dish in it. Any hydrogen sul- 
fide evolved collects in the U-tube, U, which contains pumice wet with con- 
centrated caustic potash solution. 

Since both the Mitscherlich and the Blondlot-Dusart tests give indications 
not only of white phosphorus, but also of phosphorus subsulfide, it was desirable 
to have a test to serve for the identification of white phosphorus with certainty 
even when the sulfide is also present. For this purpose, R. Schenk and E. 
Scharfff make use of the property that white phosphorus has of ionizing the 
atmosphere, a property which the sulfide does not possess. They use the 
Elster-Geitel apparatus for this purpose. For details of the test the original 
paper must be consulted. 





* A small blowpipe tip can be used here or, still better, a cylinder made by 
rolling together some platinum foil. 

+R. Scuenk and E. Scuarrr, Ber., 1906, 1522. For the detection of white 
phosphorus in the presence of hypophosphorous and arsenious acids see A. 
LecuivE, Chem. Zentr., 1912, I, 684, 


THIOSULFURIC ACID, H2S203 


This very unstable acid, in which one atom of sulfur has a positive 
valence of six and the other a negative valence of two, is soon decom- 
posed, even in dilute aqueous solution, into sulfurous acid and sulfur: 


HeS203 = H2S03+S. 


If the aqueous solution of a thiosulfate is treated with dilute hydro- 
chloric or sulfuric acid, the solution remains clear for a short time; 
but it soon becomes turbid, owing to the deposition of sulfur, which 
in this case (unlike most precipitated sulfur) appears yellow. 

The salts of thiosulfuric acid, the thiosulfates, are much more 
stable than the free acid. 


Formation of Thiosulfates 
1. By boiling sulfur with an alkali or alkaline-earth hydroxide: 
4S+60H- — 28-+8.0;-+3H.,0. 


This reaction is analogous to the action of the halogens and of phosphorus 
upon hydroxides, forming chloride and hypochlorite, phosphide (phosphine) 
and hypophosphite, etc. (cf. pp. 292 and 345). 

2. By boiling sulfites with sulfur: 


SO;-+S — 8.037. 
3. By treating alkali polysulfides with alkali sulfite in the cold: 
Na,S;+4Na,SO; =4Na,8,0,+Na:8. 
4, By the oxidation of polysulfides: 
2NaS2+302 =2Na.8.03. 


This last reaction takes place on boiling the solution of polysulfide in the 
air, or very slowly on standing. Yellow ammonium polysulfide is changed, on 
standing in the air, into ammonium thiosulfate with deposition of sulfur. 

The sulfites can be kept well in aqueous solutions, provided they are not 
subjected to the action of carbon dioxide. They are gradually decomposed by 
the latter, with separation of sulfur. 

The most important commercial thiosulfate is the sodium salt NasS.0;-5H,0, 
the well-known ‘ hypo ” of photographers. 


Solubility.—The thiosulfates of the alkalies are readily soluble 
in water, the remaining ones are difficultly soluble; many of them 
dissolve in an excess of sodium thiosulfate, forming complex ions. 

387 


388 REACTIONS OF THE ACID CONSTITUENTS 


REACTIONS IN THE WET WAY 


Use a solution of sodium thiosulfate for these reactions. 

1. Sulfuric Acid.—Both dilute and concentrated sulfuric acid 
decompose thiosulfates, with deposition of yellow sulfur. 

2. Silver Nitrate produces a white precipitate, which rapidly 
becomes yellow, brown, and finally black, owing to the formation of 
silver sulfide: 


S2037+2Agt — AgoS203; 
AgoS203 +H:2O = HeS04 +AgoS. 


Silver thiosulfate is soluble in an excess of the reagent. Diffi- 
cultly soluble Na[AgS2Os3] is at first formed, ’ 


Ag2S203+8203° — 2[AgS20s!, 


which combines with more thiosulfate, forming a soluble complex 
salt: 


2[AgS203]"+8203— — [Age(S203)3])" ~. 
But by boiling the dilute solution, silver sulfide is precipitated: 


Ago(S203)3]~ ~ > S2037 +804" +802 T +S+AgeS. 


Many other metals behave like silver, especially those of the 
hydrogen sulfide group. Thus copper, mercurous, and tin salts are 
precipitated as sulfides by boiling the acid solutions with sodium 
thiosulfate.* 

3. Barium Chloride in excess produces a white, crystalline pre- 
cipitate of barium thiosulfate,t difficultly soluble in cold water (500 
cc. of water at 18° dissolve about 1 gm. of BaS2O3), but fairly soluble 
in hot water. 

4. Strontium Chloride produces a white, crystalline precipitate, 
but only in very concentrated solutions (3.7 cc. of water at 18° dis- 
solve 1 gm. of SrS2Os). 

5. Lead Acetate precipitates white lead thiosulfate, solulite in an 
excess of the alkali thiosulfate. On boiling the solution a volumi- 
nous precipitate, consisting of lead sulfate and lead sulfide, is formed. 

6. Iodine Solution is decolorized by a thiosulfate solution: 


28203 +12 — 21°-+840¢6-. 





* Z. anorg. Chem., 28, 223 (1902). S 
{ Rubbing the sides of the test-tube hastens the formation of this precipitate. 


THIOSULFURIC ACID 389 


The iodine is reduced to iodine anions and the thiosulfate ion is oxidized 
to tetrathionate ion. 
Chlorine and bromine in excess (cf. D. 305) act quite differently 
upon thiosulfates. If chlorine (or bromine) is conducted into a solu- 
tion of sodium thiosulfate, a considerable precipitation of sulfur takes 
place, which, upon further action of the halogen, disappears: 


S203° +He0+Cle — 2Cl>+ 2HtS04-+S; 
S+4H.0+3Cle=8H*t+6Cl- +804". 


Other weak oxidizing agents act in the same way as iodine. Thus, 

7. Ferric Chloride produces, in’ solutions of sodium thiosulfate, 
at first a dark-violet coloration (perhaps ferric thiosulfate), which 
disappears after some time, leaving a colorless solution containing 
ferrous chloride and sodium tetrathionate: 


28203° + 2Fet + 33 = 2Fet t48406. 
Similarly, 
8. Cupric Salts are reduced to colorless cuprous compounds, with 
the formation of sodium tetrathionate: 


25203 + 2Cut + => Cut T4. S406 . 


The unstable cuprous sulfate immediately acts upon more thio- 
sulfate, forming sodium cuprous thiosulfate: 


Ug? *+4+280037 — [Cue(S203)2]>. 


If the colorless solution of the cuprous salt is treated with caustic potash 
solution, yellow cuprous hydroxide is in some cases immediately formed, in 
other cases only on standing or on warming. The precipitate becomes darker 
colored on being boiled. 

If the solution is acidified and boiled, black cuprous sulfide is precipitated. 

The colorless solution of the cuprous salt also gives a white (usually a light 
pink) precipitate with potassium ferrocyanide or cuprous ferrocyanide. 


9. Nascent Hydrogen (zinc and hydrochloric acid) causes the 
evolution of hydrogen sulfide. 

10. Zinc Salts produce no precipitate (difference from sulfides). 

11. Zine Sulfate and Sodium Nitroprusside produce no red colora- 
ation (difference from sulfites). 

12. Potassium Cyanide.—Boiling a solution of a thiosulfate with 
potassium cyanide and caustic soda transforms the thiosulfate into aut 
fite and the cyanide into thiocyanate: 


$203--++-CN~ — SO3"=+CNS~ 


390 REACTIONS OF THE ACID CONSTITUENTS 


On acidifying the solution with hydrochloric acid and adding 
ferric chloride, the blood-red color of ferric thiocyanate is obtained 
(difference from sulfites). 


Detection of Sulfurous and Thiosulfuric Acids in the Presence of 
Hydrogen Sulfide 


A Method of E. Votocek.*—Principle.—Alkali sulfites, sulfides and polysul- 


fides in slightly alkaline solution will decolorize fuchsin, malachite green or a 
mixture of these two dyestuffs. Ifa solution of acetaldehyde or of formalin is 
added to the decolorized solution, the color returns. Sulfhydrates, thiosulfates 
and thionates do not decolorize a solution of the above dyestuffs. 

Reagent.—Dissolve 0.025 gm. of fuchsin and 0.025 gm. of malachite green 
separately in 100 cc. portions of water. Mix three volumes of the fuchsin 
solution with one volume of the malachite green solution. 

Procedure.—It is assumed that the solution is slightly alkaline. Test the 
solution first for sulfide (monosulfide, sulfhydrate and polysulfide) by treating 
a little of it with 2 or 3 drops of sodium nitroprusside solution. A reddish- 
violet color shows the presence of the sulfide anion. If sulfide is present, treat 
the remainder of the solution with cadmium carbonate, shake vigorously, and 
allow the cadmium sulfide to settle somewhat. Filter and test a new portion 
of the filtrate with sodium nitroprusside to see if all of the sulfide has been 
removed. When all the sulfide has been removed or proved absent, treat the 
remainder of the filtrate with a drop of phenolphthalein solution and intro- 
duce carbon dioxide gas until the solution is decolorized by it. Take 2 or 3 ce. 
of this colorless solution and test it with 2 or 3 drops of the fuchsin-malachite- 
green reagent. If the color solution is decolorized, a sulfite is present. To the 
remainder of the solution add a little dilute hydrochloric acid, boil a few minutes, 
and notice whether there is any deposition of sulfur. If the solution remains 
clear, no thiosulfate is present. 


This is the best method for detecting a sulfide, a sulfite and a thiosulfate | 


in the presence of one another. 

(b) Method of Autenrieth and Windaus.}—The three acids are assumed 
to be present together in solution in the form of their alkali salts. Treat the 
fairly concentrated solution with cadmium carbonate, shake and filter off the 
excess cadmium carbonate and any cadmium sulfide which will be formed if 
a sulfide is present. Treat the filtrate with strontium nitrate solution and 
allow it to stand overnight. Filter off any strontium sulfite that may be 
formed and wash it with a little cold water. If the strontium sulfite is treated 
on the filter with dilute hydrochloric acid, sulfurous acid goes into solution, 
which can be detected by its property of decolorizing an iodine solution. In 
the filtrate from the strontium sulfite, the thiosulfate remains; it can be 


detected by acidifying with hydrochloric acid and warming, when sulfur will 


be deposited, 





* Ber., 40, 414 (1907). | 
+ Z. anal. chem., 1898, 295. For another method of detecting sulfite in the 
presence of thiosulfate, cf. F. E. Weston, Chem. Zentr., 1910, I, 379. 


. 
OU 


THIOSULFURIC ACID 391 


Solubility of Sulfites and Thiosulfates of the Alkaline Earths in Water. 


Sulfite. Thiosulfate. 
SERSOMIR I OR aah eis ake: 1 : 800 Fee 
Strontium..... relate PS Sra ean 1 : 30,000 LaF 
ESCM A rear anak eee ae 1 : 46,000 1 : 480 


REACTIONS IN THE DRY WAY 


The thiosulfites of the alkalies, on being heated out of contact 
with the air, are changed into sulfate and polysulfide, and the latter 
into sulfide and sulfur: 


4Na2Se03 — 3Na2SO4+NasSs, 
Nass ae Naes +48. 


If this reaction is performed in a closed tube, a sublimate of sul- 
fur is obtained (difference from sulfites); and the residue yields hydro- 
gen sulfide if treated with acid. 


GROUP V 


Silver Nitrate produces no precipitate in acid or neutral solutions. 
Barium Chloride, also, causes no precipitation. 


NITRIC ACID, HNO; 


Occurrence.—Nitric acid is found in the form of nitrates in small 
amounts almost everywhere in nature; thus the ammonium salt 
is found in the atmosphere and in soils; the calcium salt is found 
in old masonry; while the sodium salt is found in rainless localities, 
particularly in Chili (Chili saltpetre). 

Nitric acid is the final product of the oxidation of ammonia; it is 
found wherever nitrogenous organic substances have been subjected 
to decay, forming ammonia. With the help of microérganisms 
(Monas nitrificans, according to Winogradsky) the ammonia is changed 
first to nitrous acid, 


2NH3+302 = 2H20+2HNOdzg, 
and by further oxidation to nitric acid: 
2HNO2+02 =2HNOs. 


Properties—Pure nitric acid is a colorless liquid, with a specific 
gravity of 1.54 at 20°. At 86° it begins to boil, with decomposition, 
giving off its anhydride, which suffers further decomposition into 
nitrogen peroxide, NOz (brown fumes), and oxygen. By the constant 
loss of N2Os, the nitric acid becomes more and more dilute and the © 
boiling-point constantly rises, until at 120.5° C. it remains constant; 
when nitric acid of specific gravity 1.414 distills over, forming a 68 
per cent acid. Ifa more dilute acid is subjected to distillation, water 
is at first given off, the boiling-point constantly rising until 120.5° C. 
is reached, when a 68 per cent acid again distills unchanged. 

Red, fuming nitric acid is obtained by conducting NOe into the 
colorless, concentrated acid. In its most concentrated condition it 
possesses a specific gravity of 1.55. 

If the fuming acid is treated with water, it is colored green, and 


vapors of nitric oxide are given off, which are colored brown on com- 
392 


NITRIC ACID 393 


ing in contact with the air. The dissolved NOz (or, better, N2O,), 
being a mixed anhydride, is changed into nitric and nitrous acids, 


N204+H20 — HNOs+HNOzg, 


and the nitrous acid, owing to the heat of reaction, is partly changed 
into nitric acid, with evolution of nitric oxide, 


: 3HNO2=H20+HNO3+2NO 7 , 
and 
NO+0O=N0Oz 7 (brown vapors). 


Nitric acid is a strong oxidizing agent (cf. p. 30). It is mono- 
basic, and, next to the halogen acids, is the strongest acid (cf, p. 10). 
It forms stable salts, which are all soluble in water; but a few of them 
are changed by water into basic salts (cf. bismuth and mercuric salts), 
insoluble in water, but soluble in dilute nitric acid. 


REACTIONS IN THE WET WAY 


As nitric acid does not form insoluble salts, it cannot be detected by means 
of precipitation; its characteristic reactions depend upon its oxidizing action. 
Great care must be exercised before deciding whether this acid is present, for 
other oxidizing substances give similar (in some cases the same) reactions. 


1. Dilute Sulfuric Acid gives no reaction (difference from nitrous 
acid). 

2. Concentrated Sulfuric Acid when heated with any nitrate causes 
evolution of yellow to brown vapors of NOs, with a characteristic 
penetrating odor. 

3. Silver Nitrate and Barium Chloride cause no precipitation. 

4. Ferrous Salts are oxidized by nitric acid, which is itself reduced 
to nitric oxide, NO. | 


If the reaction takes place in the cold, the latter combines with the excess 
of ferrous salt, forming a dark-brown, very unstable compound, FeX,-NO. 
This compound is decomposed, on warming, into ferrous salt and nitric oxide 
(which escapes) the brown color disappearing. If the amount of nitric acid 
present is more than sufficient to oxidize completely the ferrous salt to ferric 
salt, a more reddish coloration is obtained owing to the formation of a salt such 
as Fes(SO4)3- 4NO. 

The oxidation of the ferrous salt takes place according to the following 


reaction, 
3Fett+NO; +4Ht — 3Fe++++2H,0+N0O T 


and is best carried out in the following manner: 
Place some of the substance to be tested for nitrate in a test-tube with 5 ce. 
~ of water, add an equal volume of concentrated sulfuric acid and cool nearly toe 


394 REACTIONS OF THE ACID CONSTITUENTS 


room temperature. by shaking the solution under running water. Carefully 
add about 5 ec. of saturated ferrous sulfate solution down the sides of the test- 
tube so that it forms a layer on top of the concentrated sulfuric acid solution. 
When a nitrate is present a brown ring is formed at the zone of contact between 
the ferrous sulfate solution and the heavier sulfuric acid solution, or if con- 
siderable nitric acid is present the whole of the ferrous sulfate solution may be 
colored. If only a little nitric acid is present, the zone may be colored pink 
owing to the formation of Fe.(SO,.);-4NO. The test is not reliable in the pres- 
ence of an iodide or a chromate. 


Nitrous acid gives the same reaction, with the difference that 
it takes place even without the addition of concentrated sulfuric acid. 

5. Indigo Solution is decolorized by warming with nitrie acid 
(as well as by other oxidizing agents). eee 

6. Potassium Iodide is not decomposed by pure, dilute nitric 
acid (difference from nitrous acid). 


If to the solution of a nitrate we add potassium iodide, a few drops of an 
acid (best acetic acid), and a little zinc, the nitric acid is reduced to nitrous 
acid, which then reacts with hydriodic acid so that the solution becomes yellow 
on account of the separation of iodine. By shaking the solution with carbon 
disulfide, the latter will be colored reddish violet, or the iodine may be 
detected by adding a little starch paste. 

The reactions which take place may be represented by the following equa- 
tions: 

Zn+NO;-+2Ht — Znt++NO,.-+H.0; 


2NO,_+2I1-+4H* — 2NO ft +2H,0+1. 


If it is desired to detect the free iodine by forming a solution in carbon 
disulfide, the latter wnder no circumstances should be added before the zine 
-has been allowed to act upon the acid solution of the nitrate and potassium 
iodide. In such a case, there will often be no separation of iodine, because 
the nascent hydrogen is used up, reducing the carbon disulfide to thioformalde- 
hyde and hydrogen sulfide, and the latter reacts with any iodine which may be 
formed, changing it back to hydriodic acid: 


CS.+2H:, = CH.S +H,§; 
H.S+2I =2HI+S8. 


7. Diphenylamine Reaction (the Lunge test *).—Reagent.—Dissolve 
0.5 gm. of diphenylamine in 100 cc. of pure, concentrated sulfuric acid 
diluted with 20 cc. of water. 


Procedure.—Place a few cubic centimeters of the diphenylamine solution 
in a test-tube and carefully cover it with the solution to be tested for nitric 
acid. If the latter is present, there is formed at the zone of contact between the 
two liquids a ring of a beautiful blue color. 





* LunGE, Z. angew. Chemie, 1894, 345. 


NITRIC ACID 395 


This very sensitive reaction is, unfortunately, also caused by nitrous, 
chloric, and selenic acids, ferric chloride, and many other oxidizing agents. 
Even fuming sulfuric acid will give it sometimes. 

In the absence of ferric and selenic salts, it is useful for detecting the presence 
of small amounts of nitrogen acids in sulfuric acid. In this case first pour the 
concentrated sulfuric acid to be tested into the test-tube and cover it with the 
specifically lighter diphenylamine solution. If 1 cc. of an acid containing 
only 3; milligram of nitrogen in a liter is used, the reaction will cause a notice- 
able coloration. If very strong fuming sulfuric acid is to be tested, dilute it first 
with concentrated sulfuric acid until it does not contain more than 20 per cent 
excess SOs. 3 


8. Brucine Reaction.—Reagent.—Dissolve 0.2 gm. of brucine in 
100 cc. of pure concentrated sulfuric acid.* 


Procedure.—Mix the solution to be tested for nitric acid with three times its 
- volume of pure, concentrated sulfuric acid, and add 1 ce. of brucine solution. 
If nitric acid is present, a red coloration quickly appears, which quickly changes 
to orange, then slowly to lemon or gold yellow, and finally becomes greenish 
yellow. Nitrous acid does not give this reaction provided it is present as 
“ nitrose,” 7.e., dissolved in concentrated sulfuric acid. Aqueous solutions of 
nitrites always yield a small amount of nitric acid when acidified with sulfuric 
acid, and consequently give the brucine reaction. 


9. Zinc in Alkaline Solution reduces nitric acid to ammonia. 
If a nitrate solution is boiled with zinc dust and an alkali, a con- 
siderable evolution of ammonia takes place: 


NO3~+4Zn+70H- — 4Zn02=+2H20+NHs } . 


Devarda’s alloy reacts much more quickly with a drop of caustic 
soda. This reaction is particularly suited for the detection of nitric 
acid in the presence of chloric acid (cf. p. 399). 


Detection of Nitric Acid in the Presence of Nitrous Acid 


With the exception of the Lunge-Lwoff method, there is no absolutely 
reliable qualitative test for the detection of traces of nitric acid in the presence 
of large amounts of nitrous acid in aqueous solution. A number of methods 
have been proposed which depend upon the destruction of the nitrous acid by 
diazotizing, but they all yield only approximate results; because, in order to 
destroy the nitrous acid, it is necessary first to set the acid itself free by the 
addition of another acid, which always causes a part of the nitrous acid to be 
changed to nitric acid; so that the latter will be detected even when no nitric 
acid was originally present. 

Large amounts of nitric acid in the presence of nitrous acid may be detected 
by the method proposed by Piccini,} in which a concentrated solution containing 





* Lunan, Z. angew. Chemie, 1894, 348. 
{ Z. anal. Chem., 19, 354. 


396 {REACTIONS OF THE ACID CONSTITUENTS 


salts of both acids is treated with a concentrated solution of urea, and then 
covered (by means of a pipette) with dilute sulfuric acid. A lively evolution of 
nitrogen and carbon dioxide ensues, which ceases in a few minutes: 


CO(NE,): +2HNO, =CO, T +38H,0+2N, 1. 
Urea 
When the evolution of gas has ceased, test the solution for nitric acid by 
means of the diphenylamine reaction. 
This reaction, however, does not take place quickly enough to prevent 
traces of nitric acid being formed according to the following equation: . 


3HNO. =H.0+HNO;+2NO 1. 


The odor of nitrous fumes is always perceptible in the escaping nitrogen,* 
which can also usually be detected by means of iodo-starch paper. The nitric 
acid which remains in the solution can be detected by means of the diphenyl- 
amine reaction. 

The nitrous acid may also be destroyed by boiling an alkaline nitrite solu-— 
tion with neutral ammonium chloride; but traces of nitric acid are always in 
evidence at the same time.f 

If the diphenylamine reaction gives a very intense coloration after the 
destruction of the nitrous acid by means of urea, the presence of nitric acid in 
the original compound is assured; but if the reaction shows that only a trace of 
nitric acid is present, it is probably due simply to small amounts of nitric acid 
formed by the destruction of the nitrous acid. 


REACTIONS IN THE DRY WAY 


By the ignition of nitrates of the alkalies, they are changed into | 
nitrites with loss of oxygen, and the latter are decomposed on stronger 
ignition into oxide: 


2KNO3=2KNO2+Ob2 fT, 
4AKNO2=2K20+4N0+02 7. 


All nitrates deflagrate on being heated on charcoal; i.e., the char- 
coal burns at the expense of the oxygen of the nitric acid, witli vivid 
scintillation. ; 





* Even at 0° and in an atmosphere of carbon dioxide. 
} By evaporating with ammonium carbonate solution the decomposition scarcely 
takes place at all. 


CHLORIC ACID, HC103 


Free chloric acid is extremely unstable, and is decomposed at 
40°, into perchloric acid with loss of chlorine and oxygen (cf. p. 170): 


3HC103 > H20+2C102 T +HC10,, 
and 
2Cl02=Cle T +202 Tf. 


The salts of the monobasic chloric acid, the chlorates, are quite stable 
and are all soluble in water. They are formed by conducting chlorine 
into hot alkali hydroxide solutions, which are not too concentrated.* 


60H +38Cle — 5CI-+ClOs3"+3H20. 


REACTION IN THE WET WAY 


1. Dilute Sulfuric Acid sets free chloric acid from chlorates, which, 
as above stated, is gradually decomposed, with loss of chlorine and 
oxygen, into perchloric acid. The solution, therefore, acts as an 
oxidizing agent, particularly on warming; it will turn iodo starch 
blue: 

ClO3-+61-+6H* > Cl-+3H20+4+3lz. 


The speed of the reaction between the free chloric acid depends upon 
the concentration of the hydrogen ion. The neutral salts do not 
- act as oxidizing agents (difference from hypochlorites). 


Test for Hypochlorite in Chlorates. 


To test for the presence of hypochlorite in an alkali chlorate, dissolve about 
2 gm. of the salt in 200 cc. of water, add 3 cc. of 10 per cent KI solution and 3 ce. 
of starch solution, but no sulfuric acid; the solution will at once turn blue if a 
trace of hypochlorite is present. As little as 0.1 mgm. of hypochlorite will 
give the test. 


2. Concentrated Sulfuric Acid decomposes all chlorates, setting 
free greenish-yellow chlorine dioxide gas, which explodes violently 
on warming: 


8KC103+3H2SO4 = 3KHS04+HC101+2C102 T +H20. 





*F, WINTELER, Z. anorg. Chem., 38, 188 (1902). 
397 


398 REACTIONS OF THE ACID CONSTITUENTS 


3. Silver Nitrate and Barium Chloride do not cause precipitation. 
4. Reducing Agents reduce chlorates to chlorides in acid, alkaline, 
and neutral solutions. 


The reduction in acid solution is effected by means of zine and dilute sul- 
furic acid, or by means of sulfurous acid: 


3Zn+ClO;-+6Ht —> 8Zn+++Cl-+3H,0; 
3S0;-+Cl0;- > Cl-+3S0--. 


The reduction in alkaline or neutral* solution is brought about by boiling 
the solution with zinc dust, or better, by means of Devarda’s alloy (cf. p. 34): 


ClO;-+3Zn+60H~ — 3Zn0.-+Cl-+3H.0. 


The residue of zinc dust (or copper, if the alloy is used) is filtered off, the 
solution acidified with nitric acid,f and silver nitrate added, when the charac- 
teristic, curdy precipitate of silver chloride is formed. 


5. Concentrated Hydrochloric Acid decomposes all chlorates, with 
evolution of chlorine: 


ClO3~--+5CI-+6H* > 3H20+3Ck f . 


This equation is not correct, for some ClO, is always mixed with the Ch. 
The following equation expresses this: 


3Cl0;-+7Cl-+10Ht — 5H.0+4Cl f +2Cl0, ft. 


The proportion of ClO, and Cl, formed is influenced by the concentration of 
the reacting substances and the temperature. 


6. Ferrous Salts.—By boiling chlorates with ferrous salts in the 
presence of dilute acid, the chlorate is quickly reduced to chloride 
(difference from perchloric acid): . 


ClO3-+6Fe* *+6H* — 3H20 +Cl"+6Fe* aS 


7. Diphenylamine reacts the same as with nitric acid. 


Detection of Hydrochloric, Nitric, and Chloric Acids in the Presence 
of One Another 


I. First, test for the presence of chlorine anions by treating a part of the 
solution with silver nitrate; a white precipitate of silver chloride shows the pres- 
ence of hydrochloric acid. Treat the remainder of the solution with silver 
sulfate solution until no further precipitation of silver chloride takes place, and 
filter off the precipitate. 


* The reaction takes place very slowly in neutral solutions. 
7 On acidifying with nitric acid a heavy precipitate of Zn(OH)2 is obtained 
which dissolves in more nitric acid. 


CHLORIC ACID 399 


Boil the filtrate with a little caustic potash (in order to expel any ammonia 
from ammonium salts which may be present), add zine dust (or Devarda’s 
alloy), and again boil; if nitric acid is present, ammonia will be given off. Filter 
off the residue, acidify the filtrate with nitric acid and treat with silver nitrate. 
If a precipitate of silver chloride is now obtained, chloric acid was originally 
present. 

II. Or, test a small part of the solution for hydrochloric acid by adding silver 
nitrate in excess, filter off the precipitate, treat the filtrate with sulfurous acid, 
and again test with silver nitrate: a precipitate of silver chloride shows the 
presence of chloric acid. Enough nitric acid must be added to dissolve any 
silver sulfite that may form. 

Test a second portion of the solution, as above, for nitric acid. 


REACTIONS IN THE DRY WAY 


On ignition, all chlorates are decomposed, forming a chloride 
with loss of oxygen. By heating on charcoal, deflagration takes 
place. 


PERCHLORIC ACID, HClO, 


Free perchloric acid is obtained by the distillation of potassium 
perchlorate with concentrated sulfuric acid. In this way the solid, 
crystalline- hydrate HClOs+H20O is obtained, and, on heating to 
110°, the anhydrous liquid acid distills off first and fumes strongly 
in the air, while the oily hydrate HClOs4+2H20 remains behind until 
the temperature reaches 203° C., when it also distills. 

The concentrated acid is very dangerous, and often explodes spon- 
taneously. In aqueous solution, however, it can be kept without 
danger. 7 

The salts of this monobasic acid, the perchlorates, are remarkably 
stable; they contain chlorine with seven positive charges and are 
isomorphous with the permanganates. The potassium salt is obtained 
from potassium chlorate. On melting the latter compound, at first 
a lively stream of oxygen is given off which, however, soon lessens 
The melt quickly becomes viscous, and consists of potassium chloride 
and potassium perchlorate, 


2KCIO3 = KCI4+ KClO4+0Oc2 7 , 


and the latter may be separated from the much more soluble potassium 
chloride by recrystallization. , 
Solubility —All perchlorates are soluble in water. 


REACTIONS IN THE WET WAY 


Perchloric acid is not attacked by concentrated sulfuric acid, 
nor reduced to chloride by zinc dust, Devarda’s alloy, sulfurous acid, 
or acid solutions of ferrous salts. . 

1. Potassium Salts precipitate the relatively insoluble, white, 
erystalline KClOs (cf. p. 81). : 

2. Silver Nitrate and Barium Chloride produce no precipitation. 


REACTIONS IN THE DRY WAY 


The perchlorates deflagrate on being heated on-charcoal; by fusing 
they lose oxygen, leaving chloride behind, which when dissolved in 
water gives all the reactions for hydrochloric acid. 

400 


PERSULFURIC ACID, H2S20g 


Pure persulfurie acid itself has never been isolated, its solution 
in sulfuric acid alone being known. It was first prepared by M. 
Marshall,* who electrolyzed fairly dilute sulfuric acid, keeping it 
very cold. During the electrolysis hydrogen ions are discharged 
at the cathode and unite to form hydrogen molecules, while HSO4 
anions are discharged at the anode and unite to form persulfuric acid: 


2H2SO4 > He T +H2820s. 


The preparation of ammonium persulfate, from which all other 
persulfates are made, is entirely analogous. 

The most important salts of persulfuric acid are those of ammo- 
nium, potassium, and barium. (NH4)2SeOg is readily soluble in 
water, and forms monoclinic crystals; K2SeOg is difficultly soluble 
in cold water, but much more soluble in hot water, from which solu- 
tion it is obtained by rapid cooling in the form of long crystals; 
BaS20g+4H20 is made by rubbing ammonium persulfate with barium 
hydroxide, and is fairly soluble in water. 


REACTIONS 


A solution of ammonium persulfate may be used. 

1. Water.—All persulfates are decomposed in aqueous solution 
(slowly in the cold, but more quickly on warming), forming sulfate, 
free sulfuric acid, and oxygen: 


2520s +2H20 — 4HSO04°- +02 fT , 
2BaS20s+ 2H20 — 2BaS0O.+2H2S04+02 t e 


A large proportion of the oxygen escapes as ozone, which can be 
detected by its odor, or by its property of turning iodo-starch paper 
blue. A dilute solution of ammonium persulfate decomposes slowly 
at 20° C., without evolution of oxygen, part of the nitrogen being. 
oxidized to nitric acid: 


8(N H4) 2820s-+6H20 = 14(NH4)HS804+2H2S04+2HNOs. 


* J. Chem. Soc., 59, 771. 





401 


402 REACTIONS OF THE ACID CONSTITUENTS 


2. Dilute Sulfuric Acid acts the same as water. 

3. Concentrated Sulfuric Acid.—If a solid: persulfate is dissolved 
in concentrated sulfuric acid at 0° C., a liquid is obtained which pos- 
sesses very strong oxidizing properties. The mixture is known as 
Caro’s acid.* For further particles concerning this acid see p. 403. — 

4. Silver Nitrate precipitates black silver peroxide: 


2Agt+ S203 +2H20 — 2HSO4 + 2H*+AgoOo. 


If, however, the concentrated solution of ammonium persulfate is — 
treated with ammonia and a very little silver nitrate, a lively evolution 
of nitrogen takes place, and the solution becomes heated to boiling. 
Silver peroxide is formed first, and oxidizes the ammonia to water, 
setting free nitrogen (catalysis).f 

5. Manganese, Cobalt, Nickel, and Lead Salts are oxidized in 
the presence of alkali to black peroxides: 


Mntt+ SeOs +40H — 2804>+H2eMn03+H20. 


In this last reaction persulfuric acid reacts exactly similarly to hydrogen 
peroxide. It may be distinguished, however, from the latter by the fact that 
it does not decolorize a solution of potassium permanganate, does not produce 
a yellow coloration with titanium sulfate, and does not react with chromie acid 
to form chromium peroxide (cf. p. 139). Ferrous salts are readily oxidized 
to ferric salts, and cerous salts are changed to yellow ceric salts by persulfates, 
but the latter are not decolorized by an excess of the persulfate, while they are — 
by hydrogen peroxide. 

Manganese and lead salts are precipitated quantitatively from neutral and 
slightly acid solutions by alkali persulfates, cobalt incompletely from neutral 
solutions and not at all from acid ones, and nickel only in the presence of alkali. 
Hydrogen peroxide produces precipitates of peroxides in all these solutions only 
in the presenceof alkali. In the presence of silver ions, which have a catalytic 
effect, manganous ions are oxidized to permanganate in ‘hot nitric acid solu- 
tions by means of alkali persulfates. 


6. Barium Chloride does not give a precipitate immediately in 
a freshly prepared cold solution of a persulfate; but, on standing 
some time, or on boiling, insoluble barium sulfate is precipitated. 





* Z. angew. Chem., 1898, 845; Ber., 34, 853 (1901); Ber., 41, 1839 (1909). 
+ Z. phys. Chem., 37, 255 (1901). 


PERSULFURIC ACID 403 


Monopersulfuric Acid (Caro’s Acid) H2:SO; 


This acid is formed by the hydrolysis of persulfuric acid, 
H.8,0;+H20 — H.S0,+H.S80,, 
and by the action of perhydrol (30 per cent hydrogen peroxide) on sulfuric acid: 
H.S0,+ H,0: — H.O+H,SO;. 


Unlike hydrogen peroxide, it does not reduce permanganates and unlike 
persulfuric acid it causes the immediate liberation of iodine from potassium 
iodide solutions. | 

It is usually assumed that hydrogen peroxide, persulfuric acid and mono- 
persulfuric acid all contain an atom of oxygen directly connected to another 
atom of oxygen; this atom of oxygen therefore has a positive and a negative 
charge residing upon it (cf. p. 33). 


GROUP VI 


Silver Nitrate produces no precipitate. 
Barium Chloride produces a white precipitate, insoluble in acids. 


SULFURIC ACID, H2SO,4 


Pure sulfuric acid at ordinary temperatures is a colorless, oily 
liquid of specific gravity 1.838; at low temperatures it is a solid. If 
the acid is subjected to distillation, it is always partially decomposed; 
heavy, white vapors of SO3 are given off first, and at 338° C. a 98 
per cent acid distills over. Ordinary commercial sulfuric acid has a 
specific gravity of 1.83-1.84, and contains 93-96 per cent H2SOx. 
It often contains lead sulfate, selenic acid, platinum, palladium, arsen- 
ious acid, the nitrogen acids, and small amounts of organic matte 
(whereby it is often colored brown) as impurities. 

Concentrated sulfuric acid is very hygroscopic, and is used, there- 
fore, for drying gases, ete. 

The anhydride of sulfuric acid, SO3, dissolves in concentrated 
sulfuric acid, forming pyrosulfuric acid, H2S207, which is solid at 
ordinary temperatures, melts at 35°, and loses SO3 at higher tempera- 
tures. It fumes strongly, and is called, therefore, fuming sulfuric 
acid. Sulfuric acid is dibasic and forms both neutral and acid salts. 

Solubility — Most: sulfates are soluble in water; calcium sulfate 
is difficultly soluble, strontium and lead sulfates are very difficultly 
soluble, while barium sulfate is practically insoluble in water. There 
are also a number of basic sulfates, Hg, Bi, Cr, which are insoluble 
in water, but are, as a rule, easily dissolved by dilute acid. 


REACTIONS IN THE WET WAY 


1. Sulfuric Acid naturally gives no reaction. 

2. Silver Nitrate causes no precipitation in dilute solutions, but 
in concentrated solutions a white crystalline precipitate is formed 
(100 ec. of water dissolve at 18° C. only 0.58 gm. of silver sulfate). 

3. Barium Chloride precipitates, from even the most dilute solu- 
tions, white barium sulfate, insoluble in acids. ; 

4, Lead Acetate precipitates white Jead sulfate, soluble in con- 

404 


SULFURIC ACID 405 


centrated sulfuric avid, ammonium acetate, and ammonium tartrate 
solutions (cf. p. 210). 
To detect the presence of SQ, in insoluble sulfates, tréat with sodium 
carbonate, whereby insoluble carbonate and soluble sodium sulfate are formed. 
Lead sulfate and calcium sulfate are easily decomposed by boiling with 
sodium carbonate solution, but barium and strontium sulfates are only incom- 
pletely decomposed by this treatment; they are much more readily attacked by 
fusing with four times as much sodium carbonate (cf. p. 109). 


5. By Nascent Hydrogen (zinc and acid) the sulfates are not 
reduced. 


REACTIONS IN THE DRY WAY 


The neutral salts of the alkalies melt with difficulty without 
being decomposed, while the acid salts of the alkalies readily give 
off water and SOz3 (cf. p. 130). | 

The sulfates of the alkaline earths and of lead do not undergo 
decomposition on ordinary ignition; the remaining sulfates are more 
or less decomposed. 

All sulfates are reduced to sodium sulfide when heated with sodium 
carbonate on charcoal; if the product is placed upon a bright silver 
coin and moistened, a black stain of silver sulfide results, e.g.:: 


(a) CaSO4+NazgCOz = CaCO3+ Na280z; 
(b) Na2SO4+2C =2COz T +Naz8; 
(c) NagS+Age+H20+0 =2Na0H+Ag.S. 


This reaction is called the Hepar reaction. 


- HYDROFLUORIC ACID, HF 


Occurrence.—Hydrofluorie acid occurs in nature only in the form 
of fluorides, of which the most important is fluorite, CaF 2, crystalliz- 
ing in the isometric system. It is also found as cryolite, Nas[AIFs], 
_ In Greenland, and in many silicates, such as tourmaline, topaz, lepido- 
lite, apophyllite, apatite, etc. 

Preparation.—Hydrofluoric acid is obtained by decomposing 
a fluoride with concentrated sulfuric acid in platinum or lead retorts: 


CaF2+He2SO4=CaSO4+2HF 7. 


Properties—Hydrofluoric acid at temperatures above 20° C. is 
a colorless gas which changes at 19.4° C. to a mobile, fuming liquid. 
The vapors possess a penetrating odor and are poisonous. When in 
contact with the skin, the acid produces painful burns. On heating 
the concentrated aqueous solution, the gas HF at first distills and then 
the 36 per cent acid. | 


Hydrofluoric acid is distinguished from all other acids by its ability to 
dissolve silicic acid, a property which is utilized technically for etching glass, 
and in the analytical laboratory for detecting fluorine and silicic acid, as well as 
for decomposing silicates. On account of this action upon glass the acid must 
be kept in platinum, wax, or hard rubber, and prepared in platinum or wick 
vessels. 

The action of hydrofluoric acid upon silicic acid takes place according & 
the equation 


SiO. 4HF =2H,0+SiF, 1, 


and the velocity of the reaction depends upon the fineness and nature of the 
material. Thus Mackintosh * found after one hour’s action of an excess _ 
of 9 per cent hydrofluoric acid upon quartz and opal powder that the quartz 
had lost only 1.56 per cent of its original weight while the opal had lost 77.28 
per cent. 

If precipitated and ignited silica is treated with strong hydrofluoric acid, 
it dissolves almost immediately with hissing and strong evolution of heat, 
while quartz powder under the same treatment is dissolved only slowly. 

Most silicates, with regard to the ease with which they are attacked by 
hydrofluoric acid, stand intermediate between the precipitated silica and quartz, 
although some silicates are attacked more difficultly than is quartz, and a few 
are only slightly acted upon. 


* Chem. News. 64, 102, 





406 


HYDROFLUORIC ACID 407 


Hydrofluoric acid is a weak, monobasic acid having, in common with other 
weak acids like carbonic acid, acetic acid, etc., the property of turning blue 
litmus red and Brazil-wood paper yellow. The aqueous solution of an alkali 
fluoride has a strong alkaline reaction. 

The property of forming very. stable complex metal-hydrofluoric acids 
is characteristic of hydrofluoric acid as of hydrocyanic acid, 


K[Ag(CN),.]:—H[AgF.], H[KF:], H[NaF.] and H[NH.F,], 
and corresponding to ferricyanic acid, H;[Fe(CN).]: 
Na;[FeF’s], Na,[AIF 6], etc: 


Unlike the complex cyanogen compounds, of which the free acids either 
do not exist at all or represent very unstable compounds, the corresponding 
fluorine acids are fairly stable. Thus hydroargentifluoric acid, H{AgF,] 
decomposes only on gently heating it, into silver fluoride and hydrofluoric 
acid, and the corresponding alkali compounds are decomposed only upon 
ignition; for this reason the latter are suitable for attacking difficultly decom- 
posable silicates, zircon and titantium minerals, etc., which are only partially 
attacked by free hydrofluoric acid. 


Solubility—The fluorides of the alkalies, of silver, aluminium, 
tin, and mercury are soluble in water, while those of the alkaline 
eniths: of lead, copper, and zinc, are insoluble, or at least very dif- 
ficultly soluble. 


REACTIONS IN THE WET WAY 


For reactions 1, 2, and 3, use powdered calcium fluoride, but for reactions 
4, 5, 6, and 7, use a solution of sodium fluoride. 


1. Dilute Sulfuric Acid causes only a slight reaction. 
2. Concentrated Sulfuric Acid reacts readily on warming, setting 
free hydrofluoric acid: 


CaF2+H2SO4=CaSO1+2HF }. 


Acid containing about 90 per cent H.SO, is most suitable for this reaction. 
Acid containing an excess of SO; is likely to cause the formation of difficultly 
volatile fluorsulfonic acid, HSO,F. 

If the reaction is performed i in a test-tube, the hydrofluoric acid will attack 
the glass, forming volatile silicon fluoride and salts of hydrofluosilicic acid; 
but the latter are decomposed by the concentrated sulfuric acid into sulfate, 
hydrofluoric acid, and silicon fluoride, 


Na,CaSi;O.,+28HF =14H,0 Tf +Na,SiF,+CaSiF,+4SiF, f , 


Soda glass 
and 


NaSiF;+H.SO, =Na,.SO,+2HF T +Sik, T ; 
CaSiF,+H.SO, =CaSO,+2HF f +SiF, ft. 


408 REACTIONS OF THE ACID CONSTITUENTS 


The silicon fluoride formed by this reaction is a colorless gas with a pene- 


trating odor, and is decomposed by water, forming gelatinous silicic acid and 
hydrofluoric acid, 


(1) SiF,+4H.0 =H,Si0.+4HF, 


Silicon fluoride, however, readily combines with hydrofluoric acid to form 
hydrofluosilicic ‘acid: 


(2) SiF.+2HF =H,[SiF 4, 


and the latter compound is not decomposed by water. The whole reaction, 
therefore, which takes place between silicon tetrafluoride and water is expressed 
by the sum of the two equations, as follews: 


38SiF.+ 4H.0 = H,Si0,+ 2H.[SiF’]. 


If, therefore, a fluoride be heated in a glass test-tube with concentrated 
sulfuric acid, and the escaping vapors allowed to act upon water by placing 
a moist glass rod in the tube, the water adhering to the rod will become 
turbid. 

Remark.—Although the above test rarely fails when relatively large 
amounts of fluoride are used, it will not be obtained in the case of certain 
minerals containing fluorine, such as topaz, tourmaline, etc. The test may 
fail, furthermore, if the fluoride is mixed with a large excess of that modifica- 
tion of silicic acid which is most readily attacked by hydrofluorie acid. Ac- 


cording to Daniel,* this is due to the formation of a stable oxyfluoride, prob- 


ably of the formula SiOF>. 
The silicon tetrafluoride at first formed combines with the excess of amor- 
phous silicic acid present, as follows, . 


SiF,+Si0, =2Si0F,, 


but this reaction will take place only very slowly, if at all, with quartz powder 
or with the silica of a silicate such as glass. 

A positive result will be obtained invariably when the tetrafluoride test is 
made in a platinum vessel with a relatively large amount of fluoride and 
comparatively little amorphous silicic acid or silicate (large amounts of quartz 
do not influence the reaction); the test will be negative, on the other hand, 
if made in platinum with no silicic acid, or, strange to say, when only quartz 
is present with the fluoride. The reason for this different behavior lies 1 in the 
difficulty with which quartz is attacked by hydrofluoric acid. 

Daniel recommends the following method for performing the test: 

Mix the substance to be tested for fluorine with about three times as much 
(by volume) ignited quartz powder, place it in a test-tube and stir it into a 
thin paste with concentrated sulfuric acid. Close the test-tube with a cork 
in which one hole has been bored and an opening cut in the side. Through 
the hole in the cork pass a glass rod blackened with asphalt paint, and on 
the bottom of the rod suspend a drop of water; push this rod down into the 
tube until it is only a distance equal to about 14 times the diameter of the 
test-tube from the paste in the bottom. Gently fine the tube and its 


*Z. anorg. Chem., 38 (1904, 299). 





‘ «A 
ee 


fh, 
i ie 


Tae Te 


HYDROFLUORIC ACID 409 


over a small flame, and if a fluoride is present a white film of H,SiO, will be 
formed in the drop of water and will be shown plainly in contrast to the black 
rod. In a tube with 1 cm. diameter, fluorine equivalent to 1 mgm. CaF, may 
be detected, while with a tube of only 0.5 cm. diameter as little as 0.1 mgm. 
CaF, will give the test. When using a tube of small diameter, it is best to add 
the sulfuric acid through a small capillary pipette to avoid wetting the sides of 

the tube. 7 : 

If the substance contains considerable amorphous silica, or when an 
oxyfluoride, such as topaz, is present, which is hard to decompose with sulfuric 
acid, the test will fail, and it is then necessary to make use of the etching test. 

3. The Etching Test.—Place the substance to be tested for fluoride in a 
platinum crucible, add some concentrated sulfuric acid, and cover the crucible 
with a watch-glass whose convex side has a thin coating of beeswax through 
which a few letters have been scratched with a pointed match. On warming 
the contents of the crucible, the glass will be etched at the places where the 
escaping gas comes in contact with it, if a fluoride was originally present. By 
covering the upper concave side of the watch-glass with a little cold water, 
the wax coating will not melt during the experiment. 

If it is desired to detect the presence of a trace of fluorine, allow the crucible 
to stand covered with the watch-glass for twelve hours and then heat for a few 
minutes. The presence of only 0.3 mgm. CaF, is sufficient to give this test, 
provided a crucible of the right size is used. 

If the fluoride contained silica (as in topaz, tourmaline, and other minerals), 
the etching test will be negative, for even if the fluorine escapes, it will be in 
the form of silicon fluoride, which does not attack glass. 

To detect small amounts of fluoride in silicates, it is necessary first to 
transform the fluorine into calcium fluoride and to subject the latter com- 
pound to the test. 

To obtain the fluorine as calcium fluoride, proceed as follows: 

Fuse the finely pulverized silicate with six to eight times as much sodium 
carbonate in a platinum crucible, and treat the melt with water after it is cold. 
A solution is thus obtained in which all of the fluorine is present as sodium 
fluoride, together with sodium silicate. Precipitate the silicic acid by adding 
considerable ammonium carbonate to the solution, warming it slightly, and 
allowing it to stand twelve hours. After filtering off the silica, evaporate 
the solution to a small volume, and add a little phenolphthalein, which will 
impart a pink color to the solution on account of its being slightly alkaline. 
Carefully add hydrochloric acid until the stirred solution becomes colorless, 
and heat to boiling, when the color will reappear. Again decolorize the solution 
with hydrochloric acid after it has becomes cold, and repeat the process until 
the solution becomes only faintly colored on boiling it. 

Now add calcium chloride solution and again boil the solution. The 
precipitate formed consists of calcium carbonate and calcium fluoride; filter it 
off, wash, dry, and ignite it in a platinum crucible. Treat the ash with dilute 
acetic acid, evaporate to dryness, triturate with water and filter off the 
undissolved calcium fluoride. After drying the precipitate and burning the 
filter, it is ready for the etching test. 


4. Silver Nitrate causes no precipitation from solutions of soluble 
fluorides. 


410 REACTIONS OF THE ACID CONSTITUENTS 


5. Barium Acetate precipitates barium fluoride soluble in an — 


excess of mineral acid and in ammonium salts. Traces of fluorine 
present as preservative in foods, liquors, ete., may be detected by 
adding a little potassium sulfate (about 0.38 gm.) to the solution, 
heating it to boiling and slowly introducing 10 ce. of 10 per cent barium 
acetate solution. The precipitate of barium sulfate and fluoride is 
then subjected to the etching test.* 

6. Calcium Chloride gives a white, slimy precipitate, difficultly 
soluble in hydrochloric and nitric acids, but almost entirely insolu- 
ble in acetic acid. On account of its slimy consistency, the precipi- 
tated calcium fluoride is extremely hard to filter; but by precipitating 


it in the presence of calcium carbonate a mixture is obtained which © 


can be readily filtered. After igniting and treating with acetic 
acid, the precipitate is changed to soluble calcium acetate and in- 
soluble calcium fluoride; it is now much denser and can be filtered 
readily. 

7. Ferric Chloride produces in concentrated solutions of alkali 
fluorides a white, crystalline precipitate corresponding to the gen- 
eral formula M;[FelF’¢]. These salts, which are analogous to cryo- 
lite, Nas[AlF], ave difficultly soluble in water, and their saturated, 
aqueous solutions do not give the iron reaction upon the addition of 
potassium thiocyanate, except after the addition of acid. These 
complex fluorides also are slightly decomposed by ammonia, forming 
a basic ferric fluoride. 


Methods for Getting Insoluble Fluorides into Solution 


(a) Calcium fluoride alone cannot be completely decomposed by fusing 
with sodium carbonate. The aqueous solution of the melt always contains 
a considerable amount of sodium fluoride, but never the total amount of the 
fluorine. If, however, the fluoride is mixed with silica or a silicate, complete 
decomposition can be effected by fusing with sodium carbonate. The silica 
decomposes calcium fluoride, forming calcium fluorsilicate and calcium silicate, 
salts which are decomposed by fusion with sodium carbonate. 

On treating the melt with water, sodium fluoride and sodium silicate go 
into solution, while the calcium is left behind on the carbonate, and can be 
dissolved by treatment with dilute hydrochloric acid. 

(b) All fluorides are decomposed by heating with concentrated sulfuric 
acid, being changed to sulfates. 





* Cf. Buarez, Chem. News, 91, 39; also WoopMAN and Taxpor, J. Am. Chem. 
Soc., 27, 1437. 


nee oe ee ae 


HYDROFLUOSILICIC ACID 411 


REACTIONS IN THE DRY WAY 


Most fluorides are unchanged by ignition. By heating them 
with silica in moist air, they are all more or less completely decom- 


posed: 
CaF2.+H20+Si02 =CaSi0;+2HF T ; 


4HF-+SiO2 =SiF4+2H20. 


The acid fluorides give off hydrofluoric acid on ignition, whereby 
the glass tube in which they are heated becomes etched. 


HYDROFLUOSILICIC ACID, H2SiFs 


As we have seen, this acid is formed by the action of silicon fluo- 
ride upon water: 


38iF4+4H.20 = 2H2[SiF6]-+HuSiOx. 


If the silicic acid is filtered off, a strongly-acid solution is obtained 
containing hydrofluosilicic acid. By evaporating the solution, the 
acid is decomposed into silicon fluoride and hydrofluoric acid, 


H[SiFs] =SiFs t +2HF ¢, 


so that hydrofluosilicic acid itself is known only in aqueous solution, 
although its salts are very stable. 

Solubility—Most silicofluorides are soluble in water; the potas- 
sium and barium salts form exceptions, being difficultly soluble in 
water and insoluble in alcohol. 


REACTIONS IN THE WET WAY 


A solution of sodium silicofluoride should be used. 

1. Dilute Sulfuric Acid causes only a very slight decomposition. 

2. Concentrated Sulfuric Acid decomposes all silicofluorides, 
evolving silicon fluoride and hydrofluoric acid: 


’ NaoSiF. + H2eSO4 = Na2SO1+Sik 4 t +2HF + : 


If the reaction is performed in a platinum crucible, the escaping 
gas will etch glass, and will cause a drop of water to become turbid. 

3. Silver Nitrate produces no precipitation. 

4. Barium Chloride gives a crystalline precipitate (1 gm. BalSiFs] 
dissolves in about 3750 cc. of water at 17°). 

5. Potassium Chloride produces, from solutions which are not 
too dilute, a gelatinous precipitate of potassium silicofluoride, which 





412 REACTIONS OF THE ACID CONSTITUENTS 


is difficultly soluble in water (1 gm. of Ko[Sik’g] dissolves in 835 ce. of 
water at 17°) and much more insoluble in an excess of potassium 
chloride or in alcohol, but soluble in ammonium chloride. 


6. Ammonia decomposes all soluble silicofluorides, with separa- 
tion of silicic acid: = 


NagSil’s +4NH40H = 2NaF+4NH4F+H,Si04. 
7. Potassium and Sodium Hydroxides react in the same way as 
ammonia, but the silicic acid remains in solution as alkali silicate. — 
REACTIONS IN THE DRY WAY 


All silicofluorides are decomposed on being heated into fluoride 
of the metal and silicon fluoride: 


K2SiF6=2KF+SiF4 T . 


The escaping gas renders a drop of water turbid, and the residue 
gives all the reactions of a fluoride 7 


> 


GROUP VII 


NON-VOLATILE ACIDS WHICH FORM SOLUBLE SALTS 
WITH THE ALKALIES 


SILICIC ACID, H4SiO, AND H2SiO3 


Occurrence.—The above acids, from which very stable salts are 
derived, are not known in the free state, as is the case with carbonic 
and sulfurous acids; although there are indeed amorphous, natural 
minerals consisting of hydrated silica with varying amounts of 
water: water opal with about 36 per cent water, ordinary opal with 
from 3 to 13 per cent water, and hyalite with about 3 per cent water: 
but none of these substances represents a compound of constant 
composition. 

The anhydride SiOz occurs in rhombohedral crystals as quartz, 
whose prismatic faces are almost always striated horizontally; and 
as tridymite, also crystallizing in the hexagonal system. The amor- 
phous silicic acid is often found mixed with the crystallized anhy- 
dride as flint, agate, chalcedony, jasper, etc. Silicic acid is, however, 
most frequently found in the form of its salts, the silicates. 

Preparation and Properties——Silicic acid can be very readily ob- 
tained pure by the hydrolysis of its fluoride, 


38iF4+4H20 = 2H2SiF's6+H4SiOng, 
or by the decomposition of alkali silicates (water-glass) with acids: 
NazSi03+2HCl=2NaCl+H2SiOsz. 


The silicic acid thus obtained forms an amorphous, gelatinous mass, 
appreciably soluble in water and acids, and readily soluble in even dilute 
solutions of caustic alkalies or alkaline carbonates. Thus freshly precipitated 
silicic acid will be readily and completely dissolved by a short digestion with 5 
per cent (or even 1 per cent) sodium carbonate solution on the water-bath. 
On being dried, silicic acid gradually loses water, and at a gentle red heat is 
changed into the form of its anhydride. According to the extent to which the 
dehydration has gone, the solubility of the silicic acid diminishes both in acids 
and in alkalies. . 

1. Air-dried silicic acid, with 16.65 per cent of water, corresponding to the 
formula 3Si0.-2H;0, is perceptibly soluble in acids, and completely dissolved 
by digestion for one-quarter to one-half an hour with 1 per cent soda solu- 
tion on the water-bath. 

413 


414 REACTIONS OF THE ACID CONSTITUENTS 


2. Silicie acid dried at 100° with 13.60 per cent of water, corresponding 
to the formula 2Si0,-H,O, is practically insuluble in acids, but can be dissolved 
by digesting for one-quarter hour with 1 per cent sodium carbonate solution 
upon the water-bath, or more readily by boiling. 

3. Silicic acid dried at 200°, with 5.66 per cent of water, corresponding 
to the formula 5SiO.-H.O, and the acid dried at 300°, with 3.40 per cent of 
water, corresponding to the formula 9S8i0.-H.O, dissolve slowly by digestion 
with i per cent sodium carbonate solution on the water-bath. 

4. The anhydride obtained by gentle ignition to a faint-red heat is only 
partly dissolved by 1 per cent or by 5 per cent sodium carbonate after half 
an hour’s digestion on the water-bath; but is dissolved after boiling for two 
hours with the sodium carbonate solution. 

5. The strongly ignited anhydride is dissolved slowly by 5 per cent sodium 
carbonate solution after repeated boiling for a long time, but is readily dis- 
solved by boiling with concentrated caustic soda or potash. 

6. The native anhydride, quartz, after being powdered in an agate mortar, 
is practically insoluble in 5 per cent sodium carbonate solution, and very diffi- 
cultly soluble in boiling caustic alkali. If it is in the form of an extremely 
fine powder, it can be dissolved by boiling with 5 per cent sodium carbonate 
solution (Lunge and Millberg). 

It follows from the above that the solubility of silicic acid (and of its 
anhydride) in alkali carbonates depends largely upon the fineness of the material. 

Silicie acid, as well as its anhydride, is soluble in aqueous hydrofluoric 
acid, forming hydrofluosilicic acid: 


Si0.++6HF =2H.0+H.SiF,. 


By evaporating this solution hydrofluoric acid is evolved; and silicon 
fluoride, with small amounts of silicic acid, is left behind. In order, then, 
to volatilize silicic acid completely by means of hydrofluoric acid, the hydro- 
lytic action of water must be prevented, which is effected by the addition of 
a little concentrated sulfuric acid. The procedure is as follows: 

Moisten the substance in a platinum crucible with a very little water, add 
not more than 3 ec. of concentrated sulfuric acid, and then the hydrofluoric 
acid. Evaporate the mixture on the water-bath, or suspend the crucible in a 
larger crucible, and heat the latter until the hydrofluoric acid is expelled, cool, 
add another portion of hydrofluoric acid and again evaporate. If a very 
large quantity of silicic acid is present it may be necessary to treat with 
hydrofluoric acid a third time. It is better to proceed in this way than to 
add a large quantity of hydrofluoric acid at one time. Finally drive off the 
sulfuric acid by heating directly over a small flame. 

The salts of silicic acid, the silicates, are exceedingly numerous, and are 
usually very stable. Many of them are so stable that they are not attacked 
by concentrated acids, while others are easily decomposed thereby. 

The different silicates are classified according to their solubility into 

A. Water-soluble silicates. 

B. Water-insoluble silicates, which are again divided into 

(a) Silicates decomposable by acids; 
(b) Silicates undecomposable by acids 


SILICIC ACID 415 


A. Water-soluble Silicates 


The silicates which are soluble in water, or ‘‘ water-glasses,” are obtained 
; £ ’ 


by fusing silica or a silicate with caustic alkali or alkali carbonate: 
Si0.+ NaCO; =Na.Si0;+CO, i} e 


1. Behavior Toward Acids.—The aqueous solution of an alkali 
silicate reacts strongly alkaline, showing that the salt is hydrolyzed 
to a marked degree: 


Si03° +2H20 — 20H + H2Si03. 


The silicic acid set free by the hydrolysis is present as hydrosole 
in the solution. By the addition of acid the alkali hydroxide is con- 
verted into salt and a part of the silicic acid is coagulated, provided-the 
solution is not too dilute. 


The precipitation is by no means quantitative; a considerable quantity 
of silicic acid remains in solution and, in fact, under some conditions all of it 
may remain dissolved in the dilute acid. If, namely, a 10 per cent water-glass 
solution is poured quickly into hydrochloric acid of specific gravity 1.1 to 
1.3, there*is no precipitation. After standing some time, however, the entire 
contents of the beaker are changed to a jelly. From 1 per cent solutions 
treated similarly with acid, no precipitate appears even after standing a year. 

The silicic acid which is precipitated upon the addition of acid is, there- 
fore, considerably soluble in dilute acids. In order to separate the silicic 
acid completely from a solution of water-glass, the hydrated acid must be 
changed into the less hydrated acid, 28i0.-H.O, by heating at 100° C. (ef. 
p. 414). For this purpose acidify the water-glass solution with hydrochloric 
acid (or nitric or sulfuric acid) and evaporate on the water-bath to complete 
dryness (the mass must no longer smell of acid). Moisten the dry residue 
with strong acid, warm slightly, dilute and filter off the silicic acid. The small 
amount of silica remaining in solution can be removed almost entirely by a 
second evaporation of the filtrate. 


2. Behavior Toward Ammonium Salts.—If a solution of water- 
glass is treated with an ammonium salt, the silicic acid will, for the 


‘most part, be precipitated as hydroxide; the precipitation is not quite 


quantitative, but more complete than is obtained by the addition 
of cold dilute acid: 


Si03~+2NH4* oe 2NH3+H2SiO3; 
NH3+H20 @ NH40OH. 


The hydroxyl ions have a marked solvent action upon the silicic 
acid. For this reason the precipitation is more complete with an 
ammonium salt of a strong acid than with that of a weak acid, which 


416 REACTIONS OF THE ACID CONSTITUENTS 


is already hydrolized to a considerable extent. Boiling off the 
ammonia helps to make the reaction complete. The use of ammonium 
carbonate, though less satisfactory than ammonium chloride, is 
necessary when it is desired to test the solution for chloride. 

Silicic acid is more completely precipitated by zinc-ammonia 
hydroxide than by ammonium carbonate, 


Si03° +[Zn(NH3)6](OH)2=20H +6NH3+ZnSiOsz, 


because the zinc silicate formed by the reaction is much more difficultly 
soluble in dilute alkaline solution than is the free silicic acid. 


The separation of silicic acid from a solution of water-glass by means of 
ammonium carbonate may be illustrated -by a common case. Many rocks 
(particularly the zircon-syenite of Norway and Greenland, many granites and 
basalts) contain small amounts of sodalite, NaCl-3NaAISi0,, a chloride silicate 
of the leucite group. In order to detect the chlorine in such a rock, the fol 
lowing process may be used: Fuse the finely powdered silicate with six 
times as much sodium carbonate in a platinum crucible, extract the product of 
the fusion with cold water and filter. The filtrate contains all of the chlorine 
as sodium chloride in the presence of sodium silicate. Treat the solution 
with ammonium carbonate, warm gently, allow to stand twelve hours, and 
then filter off the precipitated silicic acid. In order to separate the rest of 
the silicic acid, add a little zinc-ammonia hydroxide and boil the solution 
until it no longer smells of ammonia. Filter off the precipitated zine silicate 
and zinc oxide, acidify the filtrate with nitric acid and test for chlorine with 
silver nitrate. 

To prepare the zinc ammonia hydroxide dissolve pure zine in nitric acid, 
treat the solution with potassium hydroxide solution until it is neutral, and 
dissolve the filtered and washed zinc hydroxide in 6-normal ammonium 
hydroxide. 


B. Silicates Insoluble in Water 
(x) Decomposable by Acids 


A large number of native silicates are decomposed by evapora- 
tion with hydrochloric acid, the silica being deposited sometime. 


in the form of a jelly and sometimes in the form of a powdery mass. _ 


All zeolites, and a number of artificial silicates (such as Portland 
and Roman cements) belong to this class of silicates. 


To remove all of the silicic acid from these silicates, treat the finely powdered 
mineral with dilute hydrochloric acid, evaporate to dryness on the water- 
bath, moisten the mass with concentrated hydrochloric acid to convert any 
oxides or basic salts of iron, aluminium, magnesium, etc., into soluble chlorides, 
heat gently, dilute with hot water, boil and filter. The silicic acid is left on 
the filter, and the filtrate contains the metals as chlorides (ef. p. 414). 

The purity of the residual silicic acid must always be tested. For this 
purpose, place the well-washed precipitate together with the filter-paper, in 


» 


SILICIC ACID AIT 


a clean platinum crucible, held in an inclined position on a triangle, and care- 
fully burn the filter-paper. ‘Treat the residue with water, concentrated 
sulfuric acid and hydrofluoric acid as directed on p. 414 and finally remove 
the excess of sulfuric acid by cautious heating over the free flame. If the 
silicic acid were pure, nothing should remain after the evaporation of the 
sulfuric acid. Almost always a small residue of aluminium and ferric oxides 
remains, which in most cases can be neglected. If considerable residue is left, 
it should always be tested for titanic acid, barium sulfate and possibly tin 
dioxide. 


(8) Silicates Undecomposable by Acids 


Most silicates, the feldspars, micas, artificial glasses, porcelain, 
etc., belong to this class. In order to remove the silicic acid from 
such substances, they must be 

1. Fused with an alkali carbonate. 
2. Fused with lead oxide or boron trioxide, or 
3. Heated with sulfuric and hydrofluoric acids. 


Silicic acid is not soluble in water except to form colloidal solutions, and 
yields scarcely any hydrogen ions. In this respect it is a very weak acid but, 
as it is practically non-volatile, it is capable of expelling the acid from the 
salts of strong acids provided the base itself is not volatile at the temperature 
at which the salt is decomposed. The silicates, therefore, are very stable 
compounds particularly toward heat. ‘The natural silicates are partly derived 
from ortho silicic acid, H.SiO., and partly from meta-silicic acid, H.SiO;, but 
like other polybasic acids, silicic acid also forms salts which are derived from 
polysilicic acids such as H,48i;03, and Hi2SisO1. The solubility of the silicate 
depends upon two factors—the solubility of the oxide of the base and the 
proportion of silicic acid which it contains. As a general rule, the salts of 
ortho- and meta-silicic acids are more soluble than those of the polysilicic 
acids. Thus sodium and potassium ortho- and meta-silicates are soluble in 
water whereas a polysilicate may contain alkali as its principal base and yet 
be undecomposable by concentrated hydrochloric acid. Such silicates may, 
however, be decomposed by hydrofluoric acid, which causes volatilization of 
the silica as silicon tetrafluoride, by melting with a solid acid such as boric 
acid, or by treating with mineral acid in a sealed tube. 

Salts of bases of which the ignited oxides are very insoluble, such as Al,O,, 
sometimes form insoluble silicates of the ortho and meta types. 

- The effect of fusing a silicate with an alkali carbonate, or with a fusible 
oxide of some metal such as lead, is to increase the proportion of base in the 
silicate molecule. When the proportion of base-is increased, the solubility of 
the silicate is also increased provided the base is itself readily soluble in acid. 
It is not at all necessary, therefore, to get all the silicic acid in the form of 
sodium silicate or of lead silicate, by fusing with sodium carbonate or with 
lead oxide, but it is sufficient if the silicate is converted into a silicate which 
is decomposable by acid. For this reason the fusion with sodium carbonate 
or with lead oxide is often said to open up the silicate. It converts the silicate 
into the ortho or meta type and makes the silicate decomposable by acid. 
Thus after fusing with sodium carbonate, for example, it will be found that 


418 REACTIONS OF THE ACID CONSTITUENTS 


part of the sodium is converted into water-soluble silicic acid and part of it 
into a double silicate which is decomposable by acid. <A part of the sodium 
and a part of the silicate can be dissolved out of the fused mass by treatment 
with hot water. | 

1. Fusion with an Alkali Carbonate—This method is commonly used when 
it is desired to detect the presence of silicic acid and of all the bases except 
the alkalies. 

Mix the finely powdered substance with 4—6 times as much calcined sodium 
carbonate (or a mixture of equal parts of sodium and potassium carbonates, 
which melts lower than sodium carbonate alone), and fuse the mixture in a 
platinum crucible, heating carefully at first to avoid spattering from too violent 
evolution of carbon dioxide. Gradually increase the temperature until the full 
heat of the burner is reached, and continue fusing until the molten mass is 
quiet, and then heat for about a quarter of an hour over the blast lamp. 
Make a spiral by winding some platinum wire around a stirring rod, and insert 
the spiral in the melt. Cool the crucible quickly by directing a blast of cold 
air against its sides, and while the contents of the crucible are still warm, but 
not hot enough to spatter badly on the water, cover with a little water from 
the wash bottle. After a few minutes the fusion can usually be withdrawn 
with the aid of the platinum spiral. Treat the product of the fusion as described 
under (a). 

2Na.Si0;+CaCO;+2Al0(ONa)+14HCl = 


=6NaCl+CaCl,+2AICl;+2H.S8i0;+CO, T +5H.0. 


During the evaporation to dryness salts like aluminum chloride (ferric chloride, 
ete.) are subject to hydrolysis and are converted to some extent into oxide or 
basic salt, insoluble in water. Therefore, in order to separate the silicic acid 
from the salts, it is first necessary to convert such oxides or basic salts back 
into chlorides. This is accomplished by moistening the dry residue with con- 
centrated hydrochloric acid. After warming the acid with the residue for 
about ten minutes, dilute with hot water, boil and filter off the silicic acid using 
an ashless filter. 

To identify the silicic acid, place the well-washed precipitate, together with 
the filter, in a weighed platinum crucible, dry cavefully by a low flame placed 
in front of the crucible, and then ignite at as low a temperature as possible, 
with the flame now at the base of the crucible, until the carbon of the filter 
is all consumed. Then, for the first time, ignite strongly, eool somewhat, 
place in a desiccator and weigh when perfectly cool. Treat with hydrofluoric 
and sulfuric acids as described on p. 414, ignite and weigh after proper cooling. 
A difference in the weights before and after the treatment with these acids 
shows not only the presence of silica, but also the quantity of it. This quan- 
titative method is necessary for the detection of small quantities of silicic acid. 

To identify the silicic acid qualitatively, Daniel’s tetrafluoride test is satis- 
factory.* 


Daniel’s Tetrafluoride Test 


Ignite the well-washed precipitate, as described above, in a platinum cru- 
cible, then triturate in a mortar with three times as much potassium-sodium 
carbonate, and fuse the mixture in the crucible. After cooling the melt, 


* Z. anorg. Chem., 38, 299 (1904). 





SILICIC ACID 419 


- 


soften it by heating with a little water and treat with dilute sulfuric acid to 
decompose the excess of carbonate as well as the salt of silicic acid formed during 
the fusion. Heat the mixture in the crucible, by placing the latter upon a 
piece of asbestos board, and evaporate nearly to dryness, or until a thick 
jelly of silicic acid remains. After cooling, add three times as much fluorspar 
as there was original precipitate, a little magnesite and enough concentrated 
sulfuric acid to make a thin paste. After mixing the contents of the crucible 
with the aid of a stout platinum wire, place a drop of water on the inside of 
a crucible cover, which is partly painted with asphaltum, place the cover on 
the crucible and heat the contents of the latter gently. From time to time, 
raise the cover to see whether the water has become turbid. It frequently 
happens that the water becomes turbid, and then, provided a large excess of 
hydrofluoric acid is present, the turbidity disappears. For this reason the cover 
must be inspected frequently in order not to miss any temporary turbidity. 

The tetrafluoride test for silicic acid is very sensitive if the reaction is 
carried out in a very small platinum crucible. If such a crucible is not at 
hand, with a capacity of say 0.5 to 1 cc., it is better to test by the quanti- 

tative method when less than 0.01 gm. of silicic acid is present. 
| 2. Fusion with Lead Oxide or Boron Trioxide——These methods are very 
rarely used in qualitative analysis, so that it will not be necessary to describe 
them here. They play a more important part in quantitative analysis and 
will be described, therefore, in the second volume of this book. 

3. Decomposition by Hydrofluoric Acid.—This method is used principally 
when a silicate is to be examined for alkalies, titanic acid or barium. Treat 
the finely powdered silicate in a platinum dish with about 2 ce. of pure sulfuric 
acid (1 vol. concentrated acid and 2 vols. of water) and about 5 cc. of freshly 
distilled hydrofluoric acid and evaporate the mixture on the water-bath, stirring 
the mass from time to time with a thick platinum wire until it no longer smells 
of hydrofluoric acid. Add 5 cc. more of hydrofluoric acid and again evaporate, 
finally heating the dishf very carefully over the free flame, under a good hood, 
until the greater part of the sulfuric acid is expelled. The mass should not 
be ignited strongly, for a part of the sulfate may then be changed to an oxide 
insoluble in water. The sulfates of iron and aluminium, for example, are 
decomposed on ignition. After cooling, treat the mass with water, and usually 
everything will gradually go into solution. If a residue remains, test it for 
barium sulfate and titanic acid. . The solution can be used for the alkali tests, 
or for the tests for the other metals, if it is desired. 


REACTIONS IN THE DRY WAY 


If silicic acid or a silicate is heated in the salt of phosphorus bead, 
the metallic oxide will dissolve, while the silicic acid itself will be left 
as a white gelatinous mass, suspended in the bead (skeleton bead). 
This reaction, however, is, not infallible for certain silicates of the 
zeolite group dissolve in the bead without the formation of the skeleton. 


SILICON, Si. At. Wt. 28.3 


Silicon exists in two modifications, one of which is crystalline, while the 
other is amorphous. Amorphous silicon is a dark-brown powder, which can 
be oxidized by heating in the air while the crystalline modification remains 
unchanged on ignition in pure air or in oxygen, but if the air contains carbon 
dioxide, it is oxidized to silicon dioxide with deposition of carbon: 


CO.+8i =810.+C. 


Crystallized silicon is not attacked by any acid, but is readily dissolved by 
boiling with concentrated caustic alkali with evolution of hydrogen: 


Si+20H-+H,0 <> Si0;-+ 2H, tT . 


Silicon unites with many metals, forming silicides. The silicides of the 
light metals, magnesium, calcium, etc., are decomposed by dilute hydrochloric 
acid with the formation of spontaneously combustible silicon hydride: 


Mg.Si+4H+ > 2Mg+++H.si f. 


The hydride of silicon is not spontaneously combustible when pure, only 
when it is contaminated with hydrogen, as is invariably the case. 

In order to detect the presence of silicon in such a compound, treat it with 
nitric acid, which oxidizes the greater part of the silicon to silicic acid. 


Detection of Silicon in Iron and Steel 


If it is a question of detecting the presence of silicon in the different kinds 
of irons (steel, cast iron, etc.) take a large amount of material, for the amount 
of iron silicide present is usually very small. Place 5 or 10 gms. of the material 
(best in the form of borings) in a large beaker and treat with 60 cc. of 6-normal 
nitric acid. A violent reaction at once takes place with evolution of brown 
nitrous fumes. As soon as this action lessens, heat the solution to boiling, 
and continue heating until no more brown fumes are given off. Then pour 
the solution into a 200-cc. casserole and evaporate as far as possible upon the 
water-bath. Heat the residue carefully over a free flame until it is perfectly 
dry and then ignite the mass until the nitrate is completely changed to oxide, 
when no more brown fumes will be evolved. After cooling, dissolve the mass 
in about 50 cc. of concentrated hydrochloric acid, heat with constant stirring 
almost to boiling, evaporate nearly to dryness, take up in water, filter, and test 
the residue for silicic acid, by seeing whether it is volatile with sulfuric and 
hydrofluoric acids. 

In the case of cast iron, the silicic acid obtained is usually considerably 
contaminated with graphite, which can be removed by long ignition in a 
platinum crucible before treating with hydrofluoric and sulfuric acids. 

420 


SILICON 421 


Detection of Silicon in Carborundum and Metal Silicides 


Other silicides, such as carborundum, SiC, are not decomposed by nitric 
acid; they can be fused with caustic alkali in a silver crucible, 


SiC+4KOH+2H,0 =K.Si0;+ K,CO;+4H, tT ; 


and on acidifying the melt, the silicic acid separates out. 

Carborundum in the form of a fine powder is also easily decomposed by 
fusing with potassium carbonate. On removing the cover of the platinum 
crucible the blue flame of burning carbon monoxide is seen: 


3K,CO;+8iC =K,.Si0,+2K,0+4C0 f. 


The method of fusing silicides with caustic alkali is often used for getting 
metallic silicides into solution. Many copper silicon alloys are scarcely attatked 
_ by even aqua regia. If, however, they are fused with caustic alkali in a sil- 
ver crucible, potassium silicate, metallic copper, and hydrogen are formed: 


SiCu.+20H-+H,0 =Si0;-+Cu,.+2H: f . 


By treating the melt with water, the soluble potassium silicate may be separated 
from the copper. . 


PART IV. SYSTEMATIC ANALYSIS 


THE purpose of a qualitative analysis is not simply to find out 
what elements are contained in a given substance, but the aim should. 
also be to get a good idea of the relative amounts that are present. 
Manganese chloride, for example, is made from pyrolusite, and 
almost always contains traces of calcium, magnesium, nickel, cobalt, 
and iron. If the analyst should report that ‘the analyzed sub- 
‘stance consists of chlorides of calcium, magnesium, nickel, cobalt, 
iron, and manganese,” it is evident that one would get but a poor 
idea of the nature of the substance. .The report should read: ‘‘ The 
substance examined was manganese chloride, and contained traces of 
calcium, magnesium, etc., as impurities.” 

In order to be able to estimate the relative amounts of the differ- 
ent components of a substance, it is necessary to start with a known 
amount (usually } to 1 gm.) and compare the size of the precipitates 
produced. It will be impossible for the beginner to estimate the 
amount of a precipitate obtained, if he has studied the reactions of 
the elements with unknown amounts of the different substances. 
If, however, he has learned to work with a known amount of material, 
he will soon be able to judge from the size of a precipitate the 
amount of element to which it corresponds. 

It is a good plan first to work through the analysis of each group 
with a known solution containing 10 mgms. of each element and then 
it is comparatively easy to determine approximately how much of each 
element is present by the test obtained in the analysis of any unknown. 
Thus, starting with 1 gm. of the original substance, it is often con- 
venient to designate as present in small quantity when apparently 
less than 10 mgms. is found, as present in medium quantity when from 
10 to 50 mgms. is found, and as present in large quantity when distinctly 
more than 50 mgms. is present. Experiments with large classes of 
students have shown that such judgments are correct in nearly nine 
cases out of ten. It should be borne in mind, moreover, that for 
estimating small quantities of substances, qualitative tests are more 
accurate than any method of quantitative analysis. Thus all col- 
orimetric methods of quantitative analysis are really based on 
qualitative tests. The comparison of a test with one obtained using 
a known quantity of substance often gives a more exact determination 

422 


PRELIMINARY, EXAMINATION 423 


of the quantity present than a method involving weighing or titra- 
tion. This is because it is easy to prepare a solution containing say 
1 mgm. per cubic centimeter with an error of less than 5 parts per 
thousand, by dissolving 1000 times as much in a liter and thoroughly 
shaking; but it is more difficult to determine with equal accuracy 
the presence of only 1 mgm. of substance. By the process of diluting 
and taking an aliquot part, it is possible to prepare a solution contain- 
ing a very small known quantity of any soluble substance. For 
convenience, it is well to have solutions at hand containing exactly 
10 mgms. per cubic centimeter of each constituent. By taking three 
small drops of such a solution, approximately 1 mgm. of the constitu- 
ent can be obtained. 

Every analysis should be divided into three parts: 

I. The preliminary examination. 

II. The examination for the metals (cations). 
III. The examination for the negative elements (anions). 
The substance analyzed may be 

A. Solid and non-metallic. 

B. A metal or an alloy. 

C. A solution (liquid). 

D. A gas. 

The whole amount of the substance at hand should never be used 
for the first analysis, but a portion should always be reserved for 
unforeseen accidents. The portion taken for analysis should be divided 
into two parts after the preliminary examination, the first part being 
used for the tests for the electro-positive and the other part for the 
tests for the electro-negative elements. 

Before beginning an analysis, the substance should be carefully 
examined with the naked eye and with the microscope, and the results 
noted. Oftentimes the odor, color, and crystalline form suffice to 
give important clues as to the nature of the substance. 


A. THE SUBSTANCE IS SOLID AND NON-METALLIC * 
I. PRELIMINARY, EXAMINATION 


This should never be omitted, for it often shows how the subse- 
quent analysis may be considerably shortened, and in some cases 
makes the further examination unnecessary. It consists only of 
making the following few simple tests: 

1. Heating in the Closed Tube.—By a closed tube is understood 
a small glass tube about 10 cm. long and 0.5 cm. in diameter sealed at 





*See p. 452 for B and C. 


424 SYSTEMATIC ANALYSIS 


one end. Place a little of the substance in the tube so that none of 
it remains adhering to the sides, hold the tube in a nearly horizontal _ 
position and cautiously heat in the flame, noting carefully whether 
any change takes place. 


The Substance is Volatile 


(a) The substance sublimes completely without any deposition of 
water; it contains no non-volatile substance. 

The sublimate is white. The halogen compounds with ammonia, 
mercurous chloride and bromide, mercuric aminochloride, arsenic 
trioxide and arsenic pentoxide may be present. 


Arsenic pentoxide melts before being changed into the trioxide. 


The sublimate is colored— 
Gray: all oxygen compounds of mercury, cyanide of mercury, 
free iodine, and arsenic. 


Mercuric cyanide leaves a brown mass, paracyanide, which only disappears 
after long-continued heating. 


Yellow: arsenic sulfide, sulfur, mercuric iodide. 


Mercuric iodide becomes red immediately on being rubbed with a glass 
rod. 


Grayish black: mercuric sulfide. 

(b) The substance is completely volatile, with separation of water 
and gaseous products: most ammonium compounds (with the excep- 
tion of those of the halogens) and free oxalic acid. 


By very cautious heating, oxalic acid may be sublimed; it usually de- 
composes, however, into water, carbon monoxide, and carbon dioxide. 


The Substance is only Partly Volatile 


In this case gases and vapors may be evolved: 

Oxygen from peroxides, nitrates, chlorates, iodates, ete. 

Carbon dioxide from carbonates and organic substances; in the 
latter case it is usually accompanied with the separation of carbon 
and evolution of empyreumatic, combustible vapors. 

Chlorine from chlorides of platinum, gold, copper, iron, ete. 

Iodine from iodides, in the presence of oxidizing substances. 

Sulfur from many sulfides and thiosulfates. 

Arsenic from arsenites and arseniates, in the presence of carbon 
or organic substances. 


PRELIMINARY EXAMINATION 425 


Arsenites are reduced without the aid of charcoal: 
-10K;As0; =6K3As0.+-6K,0+Asi. 


Water from substances containing water of crystallization, trom 
acid salts, organic substances, or from the phosphate, borate, chromate, 
vanadate, and tungstate of ammonium. 

The water given off condenses in the cooler part of the tube and 
should be tested with litmus-paper. If it reacts alkaline, it comes 
from ammonium compounds; if acid, it results from easily decom- 
posable salts of the stronger acids. 

-Many fluorides when heated with water give off hydrofluoric 
acid, which etches the glass. 

If a sublimate is formed, make the following experiment: 

Mix a little of the substance with three times as much calcined 
sodium carbonate and heat in the closed tube. If ammonium salts are 
present, the smell of ammonia can be detected. Mercury compounds 
give a deposit of gray metal (cf. p. 203); arsenic and its oxygen com- 
pounds also usually yield the gray metal (but no globules), accom- 
panied by a garlic odor. 


The oxygen compounds of arsenic do not give the metal when heated with 
pure sodium carbonate: Commercial sodium carbonate, however, is usually 
contaminated with enough paper fibers to cause the reduction. 


2. Test the Substance in the Bead. Make a borax or sodium 
phosphate bead in the loop of a very thin platinum wire (as described 
on p. 64), introduce it with a little of the substance into the oxidizing 
flame, observe the color of the bead both when it is hot and when it is 
cold, and then heat it in the reducing flame. Borax is usually used for 
this experiment, except when it is desired to test for silicic or titanic 
acids, or when the substance is white, in which case salt of phosphorus 
is used. Only colored oxides are capable of coloring the borax bead. 


Some oxides are reduced to metal, so that the bead appears gray in the 
reducing flame (see following table). CuSO, is white when anhydrous, but 
becomes blue immediately on the addition of water. 


The following substances impart a characteristic color to the bead: 
iron, manganese, nickel, cobalt, chromium, uranium, copper (didy- 
mium, cerium, vanadium, titanium, and tungsten). 

Since the coloration varies with the temperature and with the 
amount of substance used, the results to be expected, with the neces- 
sary conditions, are summarized in the table on p. 426. The following 
abbreviations will be used: h=hot; c=cold; h—c=hot and cold; 
s.s.=slightly saturated; sat.=saturated. | 


426 


SYSTEMATIC ANALYSIS 





Wito Borax. 


Wits SALT or PHOSPHORUS. — 
































Color of 
the Bead. In the Oxidizing In the Reducing In the Oxidizing In the Reducing 
Flame. Flame. Flame. Flame. 
SiO. (without |SiO. (without | SiO. (usually | SiO. (usually 
skeleton), al-| skeleton), al-| with skeleton),| with skeleton), 
kaline earths,| kaline earths] alkalineearths| alkalineearths 
Colorless Hg, Pb, Bi,| and earths,}| and earths | and earths, 
Sb, Cd, Zn,| Mn, Di, Ce, (sat. = tur-| Mn, Di, Ce, 
Sn, Ti Cu (s.s.) bid) Cu (s.s.) 
W, Mo, Fe : 
(s.s.—c) W, Ti 
G Ag, Pb, Bi, Sb, Ag, Pd, Bi, Sb, 
sand Cd, Zn, Ni Cd, Zn, Ni 
: Fe (s.s.—h), Ag Fe (s.s.—h), Ag 
Yellow (h), Ce (h), Uj ,_.. (h), Fe (sat.— 
(or th), Vv ogre m c), Ce (h), V|Fe (h), Ti (h) 
brown) sat.), Ni (c) ery (h), U (h), Ni 
(brown) (c) (brown) 
| Cr (c), Cu (h) 
Fe (h—c), Ur, : ;1Cr (c),. U (e) 
Green Cr (c), Cu (h) Cr. V (h) ai (h), U Vv (ae Mo fey 
(c—sat.) 
Co (h—e), Co (h—c), 
Blue a Ry Co (h—c) Cu [Co the), W (6) 
Mn (h—ce), Di : 
Violet || (h—c),.and Ni any at Dl oe 
(with cobalt) ¢) 
Cu (sat.), 
opaque; when Cu as in the 
very slightly borax bead; 
Red Fe (h—sat.) saturated and}| Fe (h—-sat.) Ti and W in 
. Ce (h) with a trace of Ce (h) the presence. 











Sn, ruby red 
and transpar- 
ent. : 





of iron=blood 
red 








3. Heat a Little of the Substance upon Charcoal before the Blow- 
pipe; if deflagration takes place a nitrate, nitrite, chlorate, iodate, 
etc., may be present. 

4. Heat the Substance with Soda upon Charcoal before the Blow- 
pipe.—Mix as much of the substance as can be taken up on the end 
of a knife-blade with twice as much sodium carbonate (as described 
on p. 68), place it in a cavity on a piece of charcoal and heat in the 
reducing flame of the blowpipe. 


PRELIMINARY EXAMINATION 427 


There is obtained: 


‘As malleable button: Au, Ag, Sn, Cu, which 
can be pressed flat in an agate mortar. 

As gray metallic particles: Pt, Fe, Ni, and Co. 
Pt may be pressed flat in an agate mortar; 
Fe, Ni, and Co are magnetic and are attracted 
| by a magnet (cf. p. 66). 


(As a brittle metallic button: Sb (white in- 
crustation), Bi (yellow incrustation). The 
button may be reduced to a powder by 
grinding in an agate mortar. 

As a malleable button: Pb (yellow incrus- 
| tation). 


White, yellow when hot: Zn. 
(c) Incrustation without metal. ; Brown: Cd. 
White: As (garlic odor). 


(a) Metal without incrustation. 


(6) Metal with incrustation... . ; 





(d) White, infusible, strongly 


luminous mass........... Ca, Sr, Mg, Al, and rare earths. 


(e) Sulfur compounds are reduced to sulfides. If the melt is placed on 
a bright silver coin and moistened with water, the silver is blackened (Hepar 
reaction). 


5. Test the Substance to See whether it Imparts Any Color to the 
Non-luminous Flame.—Introduce a little of the substance on a plati- 
num wire into the base of the flame (cf. p. 62), and then into the 
fusing zone. Afterwards moisten it with dilute hydrochloric acid 
and repeat the experiment. The following indications may be ob- 
tained: 

Sodium gives a yellow monochromatic flame; a piece of sealing- 
wax or a crystal of potassium dichromate appears yellow when illumi- 
nated by this flame. 

Potassium (cesium and rubidium) gives a violet flame which is 
completely obliterated by the sodium flame. If the flame is observed 
through cobalt glass, the sodium flame disappears and the potassium 
flame appears pink. 

Lithium gives a carmine-red flame (or a red line in the spectro- 
scope). 

Strontium also gives a carmine-red flame (which the spectroscope 
shows to consist of severa! lines in the orange, and a bright line in the 
blue). 

Calcium gives a brick-red flame (in the spectroscope an orange 
and a green line are seen, both about an equal distance away from 
the sodium line). 


428 SYSTEMATIC ANALYSIS 


Barium gives a greenish-yellow flame. 


In the case of barium sulfate the green flame is either only indis- 


tinctly visible or-not at all. In order to detect barium in this case, 
heat a small portion of the substance in the upper reducing flame; 
after cooling moisten it with hydrochloric acid (odor of hydrogen 
sulfide) and again heat, when the barium flame can be easily seen. 

Thallium gives an emerald-green flame. 

If a green flame is obtained, test another portion of the substance 
for boric acid, by treating with concentrated sulfuric acid and bringing 
near the flame. A green color indicates the presence of boric acid, 
but if copper is present this test is not reliable. 


By heating the solid‘substance with potassium ethyl sulfate in a test-tube, 
boric acid is converted into B(OC:H;)s, which is volatile and burns with a 
green flame. Copper chloride does not interfere with this test. 

Lead, Arsenic, Antimony color the flame light blue, and copper 
compounds color the flame either green or blue. 


Preliminary Examination for the Electro-negative Elements 
(Anions) 


1. Dilute Sulfuric Acid (2-normal).—Treat about a gram of the 
substance in a small test-tube with dilute sulfuric acid, and note 
whether a reaction takes place in the cold or not (evolution of a gas). 

The following gases can be recognized: 

HCN from cyanides (odor) ; 

Mercuric cyanide does not liberate HCN in this test. 

H2S from soluble sulfides (odor, and blackening of lead acetate 
paper) ; 

NOz from nitrites (brown fumes) ; 

SO2 without separation of sulfur from sulfites (odor of burning 
sulfur) ; 

SOz accompanied by separation of sulfur from thiosulfates; the 
-deposited sulfur is yellow, particularly after warming; . 

COz from carbonates or cyanates (barium hydroxide solution is 
rendered turbid). i 

By boiling with dilute sulfuric acid, soluble ferro- and ferri- 
cyanides are decomposed and evolve hydrocyanic acid; acetates set 
free acetic acid; hypochlorites evolve chlorine (which also takes place 
in the cold); while the peroxides of the alkalies and alkaline earths 
are decomposed with evolution of oxygen. 


Alkali peroxides also evolve oxygen when treated with water. Cf. p. 83. 





PRELIMINARY EXAMINATION 429 


2. Concentrated Sulfuric Acid.—If the substance does not react 
with dilute sulfuric acid add 3 or 4 ce. of concentrated sulfuric acid 
and heat. If the substance reacted with dilute sulfuric acid, it will 
react violently with concentrated sulfuric acid and the gas will come 
off so quickly that it will carry small particles of the sulfuric acid with 
it, which makes the gas appear to have a penetrating odor and may 
lead to a mistaken conclusion, especially as it will also cause barium 
hydroxide solution to become turbid. 

In such a case, add dilute sulfuric acid drop by drop to a new 
portion of the substance until no further action takes place, then 
add 5 ce. of concentrated sulfuric acid and heat the mixture. 

Gases and vapors may be evolved, which are 


(a) Colorless 


HCl from chlorides, fuming in the air, with penetrating odor. 
The fumes do not cause a turbidity with water. 


AgCl and HgCl, evolve HCl very slowly; the same is true of Hg,Cl., and 
in this case SQ, is also set free. Cf. p. 288. 


SiF4 from fluorides, fuming in the air, with a penetrating odor, 
and causing a turbidity on coming in contact with water. 


SiF, is formed on account of the experiment being performed in glass. In 
platinum and in the absence of silica, HF would be evolved, which does not 
render water turbid. 


SOz, without separation of sulfur. If there was no evolution of 
sulfur dioxide on treatment of the substance with dilute sulfuric acid, 
the sulfur dioxide which now escapes must come from the sulfuric 
acid itself; a metal, sulfur, a sulfide, carbon, or non-volatile organic 
matter, such as tartaric acid, citric acid, sugar, starch, etc., must be 
present. If non-volatile organic matter is present, carbonization 
will take place on warming. 

SOz with separation of sulfur indicates the presence of a sulfo- 
cyanate, in case there was no action with dilute sulfuric acid. 

CO from oxalates and other organic substances, and cyanates. 
It is an odorless gas, which does not fume in the air and burns with 
a blue flame. | 


(b) Colored. 


Cl, a yellow gas with a suffocating odor, turns iodo-starch paper 
blue, and indicates the presence of both a chloride and an oxidizing 
substance. 


430 SYSTEMATIC ANALYSIS 


ClO2, a yellow gas, very similar to chlorine, but which explodes 
violently on being heated, indicates a chlorate. If the substance — 
deflagrates on being heated on charcoal, only a small portion of the 
substance should be used for the test with concentrated sulfuric acid; 
but if no explosion takes place on warming, more of the substance 
should be added. 

HBr from bromides has a penetrating odor, fumes in the air, and is 
always colored yellowish brown by the presence of small amounts 
of bromine. The sulfuric acid is at first colored brown in the case 
of a colorless bromide, but becomes colorless on being boiled. 

CrO2Clz, brown (similar to bromine), results from the presence of 
a chloride and chromic acid. 

Iz, violet. In the case of a colorless iodide, the sulfuric acid is 
at first colored brown by small amounts of iodide, or gray, solid iodine 
is deposited if considerable iodide is present, which volatilizes on 
warming, forming violet vapors. If considerable iodide is used for 
this test, the sulfuric acid is reduced to SOz, or even H2S (ef. p. 301). 

Mn207, violet, is formed from permanganic acid, and is devom- 
posed with scintillation, often exploding, on being warmed. 

NOz, brown, with a.penetrating odor, comes from nitrates. 

After the preceding tests have been made, the next step is the 


Solution of the Substance 


As solvents the following are used: 

1. Water; 

2. Hydrochloric acid; 

3. Nitric acid; 

4. Aqua regia. 

In the majority of cases the first three solvents suffice, aqua regia 
being seldom necessary, as will be seen from the following table: 


SUBSTANCES SOLUBLE IN WATER 


Of Group I (p. 284) the following are soluble: 

1. Chlorides.—All except AgCl, CusCle, HgeCle, PtCle, AuCl, 
BiOCl1, SbOCl1, Mg2OCle. PbCle and TIC! are difficultly soluble. 

2. Bromides.—The same as the chlorides. 

3. Iodides.—All except AgI, HgeI2, Hgle, Cusl2, PdIz, Til; PbI2 
is very difficultly soluble. 

4, Cyanides.—Only the cyanides of the alkalies, alkaline earths, 
and mercury. 


SOLUTION OF THE SUBSTANCE 431 


5. Ferrocyanides.—Only those of the alkalies and alkaline earths. 

6. Ferricyanides.—Same as the ferrocyanides. 

7. Cobalticyanides.—Only those of the alkalies, alkaline earths, 
and. the ferric, mercuric, and lead salts. 

8. Thiocyanates.—Those of the alkalies, alkaline earths, iron, 
cupric copper, and mercuric mercury. 

9. Hypochlorates.—All. 


Of Group II (p. 284) the following are soluble: 

10. Nitrites—All. Silver nitrite is difficultly soluble. 

11. Acetates.—Silver and mercurous acetates and certain basic 
acetates are difficultly soluble. 

12. Cyanates.—Those of the alkalies, alkaline earths, and most 
of the remaining ones. Silver and lead cyanates are insoluble. 

13. Sulfides—Only those of the alkalies and alkalines earths. 
CaS is difficultly soluble. 
14. Hypophosphites.—All. 


Of Group III (p. 284) the following are soluble: 

15. Sulfites —Those of the alkalies, and the bisulfites of the alkaline 
earths. 

16. Carbonates.—Those of the alkalies, and the bicarbonates of 
Ca, Sr, Ba, Mg, Fe, Mn. 

17. Oxalates.—Those of the alkalies; the remainder are diffi- 
cultly soluble or insoluble. Most oxalates, however, with the exception 
of Ba, Ca, and Sr oxalates, form soluble complex salts with alkali 
oxalates. | 

18. Iodates.—Only those of the alkalies. 

19. Borates.—Those of the alkalies. The remaining borates are all 
difficultly soluble in water, but soluble in ammonium chloride as a rule. 

20. Molybdates.—Only those of the alkalies. 

21. Selenites.—Those of the alkalies are readily soluble, the re- 
maining ones are difficultly soluble. 

22. Selenates.—All except the barium and lead salts. 

23. Tellurites.—Only those of the alkalies. 

24. Tellurates.—Only those of the alkalies. 

25. Tartrates.—The normal tartrates of the alkalies, and lithium 
and sodium bitartrates. The remaining tartrates are insoluble in 
water, but are usually soluble in an excess of alkali tartrate solution, 
forming complex salts. 

26. Citrates.—Only those of the alkalies are readily soluble in 
water. The insoluble citrates usually dissolve in an excess of alkali 
citrate solution. 















































but more soluble in acids; 


stances, small letters for the unimportant ones. 


i,-slightly soluble in water and slightly soluble in acids; 
The salts referred to are generally consideféd™ 





432 SYSTEMATIC ANALYSIS 
SOLUBILITY 
ie: 2, 34 se | 8, | ac let at] + i af 
Os } Wiw;wiw|] Wi wiws|wiwl|wiwil wl wa 
ak aN } Wiwiwi|alatftalatatsltatatataia 
aS \ Wiwiwia a a|jajsj—|—|—latstatla 
Borate, BC;=..| W | W | w |w-a| a aj/al|as}a|a|s |) eae 
* Bromide, Br~..|W|Wi]wjf})wt{|wiifiwtswti]w {\|wki| w| wl] wlw 
Carbonate, Viwiwiw|a|alajal—|—|al|a]|a]a 
Chlorate, ClO;-|} W| w|wi|w!]WwWfiwtiw wi|wi|wl|wl|w 
Chloride, ClI-...; W| W|W|W!] W|W|W well W| W| WwW) WwW | 
eerie } Wi WwWiwi wi w-alwe-a} a ==). | W | 
Cyanide, CN.... W | W| w] wiew | wiw-al —|a|Ada J ai} ai 
MeCN ee } WiwWiwiwiw * wil—|—l|a I I I 
¥ [Fe ie Wiwiwiw|w w-a| —|—|A-I| a jor 
Fluoride, F~ w | w | W | a-1| A-I | w-al| w-a| a |w-a| a | w-a| w-a| w-a 
Hydroxide, \)wilwilwi|alwalwi/wi/a|a/al|alala 
Iodide, I~..... wiwiwiw| WwW |wi|wiw w|Wwiwi]w 
Nitrate, NO;~ | W Ww Wiw|wiwiwiwiwiwiwiwiw 
Oxalate,C.0.-|W/}W;}W|]aj|A | alaslai|w-al a |w-al ala 
Oxide, O..... wilwf{]—]A{W-A| W/] WiA-T;A-T}] A} A} AIL A 
eae lwiwi/wi/alalalaAlaAlal/alAalAala 
Silicate, SiOZ--| W | W| —]| a a ajajA-I}a/aljafatsa 
Sulfate, SO.-...;W | W|W|WyW-l| I] I | W/W-ll WwW} ws) wi w 
Sulfide, S=..... Wi|wiwi!]a{f{w-A|Wiwisajfjai]/A;}|A!]AIA 
Peewee, } Wiwiwiwis]wfwl|w wiwiwi{|wiw 
Shree ke } W|W!]W {veal a a | a w | a |w-al a’| W 
W or w, soluble in water; A ora, insoluble in water, soluble in HCl, HNOs or aqua regia; 


A-I 


*Based on the table in Fresenius- 


wf 















































SOLUTION OF THE SUBSTANCE 433 
‘TABLE * 
ies seek, ea fc: Oe ea Bee 

ed aldlGlelegibealee 

wi|Wi w | W{|w-al w |W] w |] wi wl] w | — |—|— eee 
Sie e@ 18 | a f ala} a | a |—|;—)—{—|— ee 
ala} ajaj|jla|ai4|ala {a |j|—|—|—s|—|— wer. 
ajal| a aij—|— /]a} a _|w-a} a |] — | — |—|— |Borate, BO= 
wiw) i |w-il ai} w | w/! w-a| W |.w|w-a|w-a| w | w |Bromide, Br~ 
Al—| a |Alal]aJ{A}] aja |—|—|—|—]/— { ae 
wiw!| wiwiw|wiiw| w]wf|wi|— |] — |—|— (Chlorate, ClO;— 
W I |W-I| A-I[} W |W|W-A| W W |W-A| W | W |Chloride, Cl— 
—|w| a |A-I| a | w-ajw!] a SLE ep ae = Ta ge es 
a-ij|—j; I a!—|} Wia}j|— ]aj—j|—l— w |Cyanide, CN 
Boe i ee) | fj) i pedi) | — 3 (TFeCN) 
E}t] i jaj—j;—j1| i faijil i |—|-|—|{ Pea 
w-alw| a | @| wj|wajla| w |w-alw| w | w |—|—|Fluoride, F~ 
Aja|—|a|—|—|a| a] a l—|a| a |—|—|{Bycronde, 
Wiw; I |wW-lI| A} A |W| a | W/W w {|w-a| a | i |Iodide, I- 
Wiw; W/|Wi|w Ww |W! W-A| W }—| — | — |—| w Nitrate, NO;- 
og 3) EE oe a a ala Sia when tt Oxalate, C,0.= 
Aj}A} A j|A|{A{A {A} A | A{AJA-I] A |—] a |Oxide, OF 

A Al 2 |} Aj}aya {}A} a+a|a| A fw-al—|— Rape ey 
aja}—}|ad—}|—]/a|j— }]a |—| — | — |—|— Sikcate, Si0,- 
W/W W-A}|A-I| W | W-A|W| w-a | W |W} — | — |—| w Sulfate, SO.= 
Bee See ACE ASIA) OA (AIA A |A| A |Sulfide, S— 
wiw| I aj};Al|W |-a}| — |w-al— —=(—1 & Bors ees 
w-a|}W| a & |w-al a@ | w| a |w-ala|—|a |—|— Ree 








I or i, insoluble in water and in HCl, HNOs or aqua regia; 
or a-1, insoluble in water, slightly soluble in acids. 


Capitals 


as normal. Acid and basic salts are omitted in this table. 
Wells Qualitative Analysis. 


-A or w-a, slightly soluble in water 


are teed TOF th 


or the more important sub- 


434 SYSTEMATIC ANALYSIS 4 


27. Pyrophosphates.—Only those of the alkalies. 
28. Metaphosphates.—Only those of the alkalies. 


Of Group IV (p. 284) the following are soluble: 

29. Phosphates.—Only those of the alkalies. 

30. Arsenites.—Only those of, the alkalies. 

31. Arseniates.—Only those of the alkalies. 

32. Thiosulfates—Almost all are soluble, though the silver and 
barium salts are difficultly soluble. 

33. Chromates.—Those of the alkalies, Ca, Sr, Mg, Zn, Mn, Fe, 
and Cu are soluble, the others are difficultly soluble or insoluble. 

34. Vanadates.—The orthovanadates are unstable; the pyro-, meta-, 
and polyvanadates are soluble in water, as a rule. The lead and 
mercurous salts are insoluble, also the vanadates of the iron group. 

35. Periodates.—All more or less soluble in water, except sliver 
periodate, which is insoluble. 


Of Group V (p. 285), the following are soluble: 

36. Nitrates.—All except a few basic salts. 

37. Chlorates.—All. 

38. Perchlorates.—All. 

39. Manganates and Permanganates.—All. 

Of Group VI (p. 285), the following are soluble: 

40. Sulfates.—All except the Ca, Ba, Sr, and Pb salts, and a few 
basic sulfates. 

41. Fluorides.—Those of the alkalies, silver, ae mercury; the 
remaining fluorides are difficultly soluble or insoluble in water. 

Of Group VII (p. 285), the following are soluble: 

42. Silicates.—Only those of the alkalies. 

43. Tungstates.—Only those of the alkalies. 


Of the salts insoluble in water, all dissolve in acid (hydrochloric or 
nitric) except AgCl, AgBr, AgI, AgCN, AuCl, PtCle, BaSO4, SrSOa, 
PbSO4, HgS, Prussian blue, CaF2, SnSe (mosaic gold), SiOz, many 
silicates, fused PbCrOxz, and the strongly ignited oxides: AleOs, CreOz, 
TiO2, SnO2z, SbeO3.* TiOes, SnOe, and Sb2O3 can be dissolved by 
long continued boiling with concentrated hydrochloric acid. 

Of the salts insoluble in acids, the following dissolve in aqua regia: 
PtCle, AuCl, HgS, SbeOzs, SnSe, and Prussian blue (after long treat- 
ment). 

The following substances are not dissolved by aqua regia: AgCl, 





* The oxides of antimony are changed to Sb.O, after long ignition in the 
alr. 


SOLUTION OF THE SUBSTANCE 435 


AgBr, AgI, AgCN, BaSO., SrSO4, PbSOs, CaF2,* fused PbCrO,, 
AlzO3z, CreO3, native TiO2 (rutile, anatase, brookite), native SnO2g 
(cassiterite, tinstone), SiOz, Si, many silicates, C, carborundum, and 
strongly ignited iridium (rhodium, ruthenium, and osmium). 

In order to bring such substances in solution it is necessary to 
subject them to a special treatment. The process to be chosen depends 
largely upon the nature of the insoluble substance, so that a few general 
tests are necessary before going farther. Very often the preliminary 
examination will have been sufficient, but it is always well to perform 
the following simple experiments: 

1. Heat a small portion of the residue insoluble in acids on the 
charcoal stick to see whether a metallic button can be produced. 

(a) No metallic button is produced. ‘The absence of silver, lead, 
and tin is thereby assured. 

(b) A metallic button is formed. The button is flattened in an 
agate mortar, and its solubility in acids is tested. 

(a) The metal dissolves in nitric acid forming a clear solution, 
showing the absence of tin. Add a little hydrochloric acid to the nitric 
acid solution; a curdy precipitate is formed if the metal is silver, con- 
sisting of silver chloride, insoluble in water, but soluble in ammonia. 

If the nitric acid solution becomes turbid on the addition of sul- 
furic acid, lead is present. 

(8) The metal does not dissolve in nitric acid forming a clear 
solution, but leaves a white, insoluble powder: metastannic acid. 
Treat a new button with concentrated hydrochloric acid, when it will 
completely dissolve if silver is absent. Mercuric chloride produces 
a white precipitate of mercurous chloride in the hydrochloric acid 
solution: tin is present. 

2. Heat a second portion of the insoluble residue in a small test- 
tube with concentrated sulfuric acid and test to see whether the esca- 
ping gas renders a drop of water turbid. 

A turbidity shows the presence of an insoluble fluoride (CaF 2). 

3. Heat another portion of the residue (with the help of a platinum 
_ wire) in the upper reducing flame of the gas-burner, allow to cool in 
the inner mantle, moisten with dilute hydrochloric acid, and notice 
whether the odor of hydrogen sulfide can be detected. Then test 
to see whether it will now impart a characteristic coloration to the 
flame. The presence of a sulfate is betrayed by the odor of hydro- 
gen sulfide, and the flame test shows whether barium alone or a mix- 
ture of barium, calcium, and strontium is present. 


* Calcium fluoride will be dissolved by the long continued action of aqua regia. 


436 SYSTEMATIC ANALYSIS 


4. Test another portion of the residue in the salt of phosphorus 
bead; silicic acid or a silicate usually gives a skeleton bead (cf: p. 
419). 

As the skeleton bead is not always obtained even when silica is 
present, a further test for silicic acid is often necessary (cf. p. 418). 

5. Now heat the salt of phosphorus bead in the reducing flame 
to test for the presence of titanium, which causes the bead to become 
violet. The violet color appears more quickly on the addition of a 
little piece of tin-foil. If iron is present at the same time, as is always 
true in the case of rutile, the bead is colored brownish red in the redu- 
cing flame. 

6. The presence of chromium is often detected by the green color 
of the residue. In the case of chromite (gray or black residue) fuse 
some of the substance with sodium carbonate and potassium nitrate in 
the loop of a platinum wire (cf. p. 142), when a yellow melt is 
obtained if chromium is present, which (after being dissolved in 
water and acidified with acetic acid) yields, with silver ee 
reddish-brown precipitate of silver chromate. 

7. If the residue is gray or black, it may also consist of carbon. 
Heat a small portion upon a piece of platinum-foil; if carbon is present, 
the mass will glow and, if it does not burn completely, a lighter- 
colored ash will be obtained. In doubtful cases melt a little potas- 
sium chlorate in a test-tube, and add a little of the insoluble residue, 
when a distinct glowing or a little explosion will take place if carbon 
is present. It is necessary to avoid the addition of shreds of filter- 
paper in this test.. _ 

8. Silicon and Silicides (carborundum, etc.) are seldom met with, 
and show the greatest stability toward the above-mentioned reagents. 
By fusing with caustic alkali in a silver crucible, however, they are 
readily decomposed with evolution of hydrogen (cf. p. 421). 

After dissolving the melt in water and acidifying, gelatinous silicic 
acid separates out, particularly after evaporation. 


MetuHops FOR GETTING SUBSTANCES INTO SOLUTION WHICH ARE ~ 


INSOLUBLE IN ALL ACIDS 


1. Insoluble Halogen Compounds (the silver compounds alone 
come into consideration) can be brought into solution by melting 
the mass, adding a little dilute sulfuric acid and a piece of zine 
so that it comes in contact with both the acid and the insoluble 
substance. After a while pour off the acid; it contains the halogen 
acid in the presence of zine sulfate, and should be kept for the sub- 





ss 


Te eee eh Te ee a ee ee a ee 


a aan 
TS SS he ee 


SOLUTION OF THE SUBSTANCE 437 


sequent tests for acids, cf. p. 306. The residue consists of metallic 
silver. Wash it with water, dissolve in dilute nitric acid, filter and 
test the solution for silver with hydrochloric acid. 

2. Insoluble Sulfates of the Alkaline Earths are brought into solu- 
tion by fusing in a platinum crucible with four to five times as much 
calcined sodium carbonate, or with a mixture of equal parts of sodium 
and potassium carbonates. Mix the finely powdered substance 
in the crucible with the sodium carbonate, cover the mixture with a 
thin layer of more carbonate, place the lid on the crucible and heat 
at first gently over a small flame in order to drive off the moisture 
which the carbonate always contains, and then raise the temperature 
until the mass fuses to a thin liquid; maintain this temperature for 
about fifteen minutes. Remove the fused mass from the crucible 
as directed on p. 418. Heat with a little water on the water bath 
until the fused mass disintegrates, and no more hard lumps can be 
felt with a glass rod, then filter. The filtrate will contain the sulfate 
as sodium sulfate, and the residue will consist of carbonates of the 
alkaline earths. Wash it with a 5 per cent sodium carbonate solution 
until no more sulfuric acid can be detected in the filtrate, and then 
wash with hot water until the wash-water no longer reacts alkaline 
(cf. p. 109). Dissolve the residue in nitric acid, and analyze as 
described on p. 110. 

3. Lead Sulfate may be boiled with a concentrated sodium car- 
bonate solution, which forms insoluble basic lead carbonate and 
soluble sodium sulfate; with caustic soda, which forms soluble lead 
plumbite and sodium sulfate; or with ammonium acetate (cf. p. 209). 
Calcium sulfate is also decomposed completely by boiling with soda 
solution, as is strontium sulfate (though less readily), but barium 
sulfate is incompletely decomposed. 

4. Silicic Acid and Silicates should be fused with sodium ae 
ate, exactly as described on p. 418. 

5. Metastannic Acid, as obtained by the oxidation of tin with 
nitric acid, is readily dissolved by boiling with a little concentrated 
hydrochloric acid, and then treating with considerable cold water 
(ef. p. 257). 

Tin dioxide, as it occurs in nature (tinstone), as well as the strongly 
ignited metastannic acid, cannot be brought into solution in this 
way. One of the methods mentioned on p. 260 (usually the 
sodium carbonate and sulfur method) must be used. 

6. Insoluble Fluorides are first heated with concentrated sulfuric 
acid, and the sulfate formed is brought into solution by the method 
described under 2. 


438 SYSTEMATIC ANALYSIS 


7. Titanium: Dioxide is fused with potassium pyrosulfate in a 
platinum crucible (cf. pp. 180 and 158); or it is fused with sodium 
carbonate, the melt treated with cold water, and the residue dissolved 
in hydrochloric acid (ef. p. 157). Heating for some time with con- 
centrated sulfuric acid will dissolve pure titanium dioxide. When 
cold the solution may be diluted. 

Fusion with potassium pyrosulfate is also suitable for decomposing 
native aluminium oxide (corundum). 

8. Chromium Trioxide and Chromite are fused with sodium ear- 
bonate and a little potassium nitrate in a platinum crucible or with 
sodium peroxide in a nickel or iron crucible (cf. p. 133). 

9. The Insoluble Complex Cyanides are completely decom 
by boiling with caustic soda in a porcelain dish. 


After boiling with the alkali, dilute with water and filter. The filtrate 
will contain the acid in the form of its sodium salt; and, in some cases, may 
also contain aluminium and zinc. Saturate the filtrate with carbon dioxide, 
boil and filter off any precipitate (Al(OH); or ZnCO;); dissolve this precipitate 
in hydrochloric acid and test for zinc and aluminium. Acidify the alkaline 
filtrate obtained above with hydrochloric acid and test for ferrocyanic and 
ferricyanic acids according to pp. 318 and 321. 


The soluble complex cyanides are decomposed before the analysis by 


heating them with concentrated sulfuric acid (cf. p. 153), 


REACTIONS THAT ACCOMPANY THE DIsSSOLVING PROCESS 


When a substance is dissolved, whether in water or in acids, 


phenomena are often observed which may be of great importance 
2s concerns the subsequent analysis. Moreover, the color, reaction 
of the solution towards indicators, or the evolution of gases will lead to 
important conclusions. First, test the substance with regard to its solu- 
bility in water, by taking about 0.5 gm. of the fine powder, adding a 
little cold water, and noting whether any bubbles of gas are given off. 

A gas is evolved when there are present: 

(a) Peroxides of the Alkalies or Alkaline Earths, which are partly 
decomposed into hydroxide and oxygen: 

2Na202+2H20 =4Na0H+02 t ; 
2Ba02-+2H20 = 2Ba(OH)2+0z2 1. 

Barium peroxide is decomposed in this way only by heating the water. 

Test the escaping gas for oxygen by means of a glowing splinter. 

In the alkaline solution (red litmus is changed to blue) some unde- 
composed peroxide will still be found. 


Dilute the solution with considerable water, cool, and carefully acidify with 
sulfuric acid, add a little ether, some potassium dichromate solution, and 


: ’ 
(a a, 


SOLUTION OF THE SUBSTANCE 439 


shake the mixture. If a peroxide is present, the upper ether layer will now 
be colored blue. A better method for detecting the hydrogen peroxide, formed 
by the action of the sulfuric acid upon the peroxide, consists in adding a few 
drops of titanium sulfate solution; a distinct yellow color will be noticed if 
- only traces of hydrogen peroxide are present (cf. pp. 84 and 159). 


(b) Carbides of the Alkaline Earths (calcium carbide). 
These are decomposed into acetylene (which has a peculiar odor, 
and burns with a luminous flame) and calcium hydroxide: 


CaC2+2H20 =Ca(OH)2+Ce2He t - 
(c) Nitrides of the Alkaline Earths (magnesium nitride). 


Magnesium nitride is decomposed by water into magnesium 
hydroxide and ammonia: 


Mg3N2+6HOH =3Mg(OH)2+2NHs3 ff. 


| If considerable water is added, there is no gas evolution, because 

the ammonia will be absorbed by the water; ‘but on boiling the solu- 
tion, ammonia will be given off, which can be readily recognized by 
its odor. 

' (d) Phosphides of the Alkalies and Alkaline Earths.—These are 
decomposed by water, setting free spontaneously combustible phos- 
phine: 

CazP2+4H20 = Pe2H4 tT +2Ca(OH)>. 


Very small quantities of the phosphide can be recognized by the 
characteristic garlic odor. 

(e) Many Chlorides, Bromides, and Iodides of the Negative Ele- 
ments, e.g., PCl3, PCls, etc., are decomposed into the halogen hydride 
and the oxygen acid of the negative element: 


PCl5+4H20 =5HCI+H3POs,. 


(f) A few Sulfides which are Seldom Met with (Mg§, AloSs, etc.).— 
These are decomposed by water with loss of hydrogen sulfide, which 
can be detected by its odor, and by its blackening lead acetate paper: 


MgS+2H.0=Mg(OH)2+H28 } . 


After any reaction caused by the first addition of water is over, 
add about 10 to 15 ec. more, heat the water to boiling and then 
allow it to cool. 

If the substance dissolves completely, forming a clear solution, 
it is evident that it is unnecessary to test for any insoluble substances 
in the subsequent analysis. 


440 SYSTEMATIC ANALYSIS 

If a residue remains, it is possible that a part of the substance — 
has dissolved in the water. To determine whether this is the case, 
decant the liquid through a filter and carefully evaporate a little of 
the filtrate to dryness on platinum-foil (or a watch-glass). If the foil — 
is heated too hot, volatile compounds may escape unnoticed. If a a 
residue remains after evaporation, it is evident that a part of the — 
original substance is soluble in water. Then treat the original residue — 
several times with small amounts of water, and analyze the aqueous — 
extract thus obtained by itself. Treat the part remaining undissolved 
with acid, using hydrochloric acid unless the preliminary examina- 
tion has shown the presence of either lead or silver, when nitric 
acid should be used. 

Treat the residue with 0.5—1 cc. of 12-normal acid (notice whethael 
there is any evolution of.a gas), heat gently, and then dilute with — | 
water, to dissolve any chlorides insoluble in hydrochloric acid. It — 
must be remembered, however, that bismuth and antimony salts — 
form insoluble basic chlorides on dilution with water, so that too — 
much water should not be added. : 

If a residue remains after treatment with acid, bring it into solu- : 
tion by one of the methods described on pp. 436-8. 


II. EXAMINATION FOR THE METALs (CaTIONS) a 


TaBLE XI.—GENERAL SCHEME FOR SEPARATING THE METALS INTO GROUPS ~ _ 


~*~ 








a 


Solution may contain all the common basic constituents. Add HCl in slight” % 
excess. (1) . 





Precipitate: | Filtrate: Groups II, II, IV and V. Saturate with HS. (8). 
Group I. F 
Examine asout-| Precipitate: | Filtrate: Groups III, IV and V. Test for “4 
lined in Table Group II. phosphoric acid. If found present modify the q 
X, p. 282. (2) | Examine as| following procedure as indicated in Table XII. 
outlined in| Add NH,OH and (NHs4) Ss. (5) ; 











Table VII, p. 
272. (4) Precipitate: | Filtrate: Groups IV and V: 
Group III. Add (NH 4)2CO3. (7) J 
If phosphate is “- 
absent exam-|  Precipitate: Filtrate: 
ine as outlined | Group IV. Group V. 


“7 


in Table V,p.| Examine  as| Examine as 
189 orinTable | outlined in| outlined in 
VI, p.192. If | Table III, p.| Table I, p. 
phosphate is| 111 or in Ta-| 97 or in Ta- 
present exam-| ble IV,p.113.| ble II, p. 99. 
ine by Table} (8) (9) Fal 
XII, p. 445. . 
(6) . . 


*i 

















EXAMINATION FOR THE METALS 441 


- PROCEDURE 


1. Add HCl as directed on p. 283. If the original substance was com- 
pletely soluble in dilute HCl, it is evident that no silver or mercurous salt is 
present. Often, when lead is present, the solution is clear while hot, but lead 
chloride is deposited as the solution cools. It is usually best to filter off 
such a precipitate, but it will be changed to less soluble lead sulfide when H.S 
is introduced to precipitate the second group. If the original solution is 
alkaline to phenolphthalein- or to litmus, a precipitate may form when none 
of the metals of the first group is present. Thus a solution of sodium silicate 
gives a white, gelatinous precipitate of silicic acid, a solution of an alkali tung- 
state gives a precipitate of tungstic acid, a solution of a thio salt of arsenic, 
antimony or tin gives a colored sulfide precipitate and the solution of a complex 
cyanide may form a precipitate when neutralized. These precipitates, however, 
are not likely to be mistaken for a chloride of silver, lead or mercury. If-a silicate 
is present, it is absolutely necessary, to remove the silicic acid at the start 
by the method given on p. 416 and tungstic acid may be removed in exactly 
the same way. If a thio salt is present, examine the precipitate according to 
Table IX, p. 275, and test for alkalies and alkaline earths according to Table 
Il], p. 111 and Table I, p. 97. If the original solution is alkaline, it is 
necessary to test for iron and aluminium only when the solution contains non- 
volatile organic matter which prevents the precipitation of these elements by 
hydroxyl ions. The addition of HCl may cause the precipitation of BiOCl 
and dilution may cause the precipitation of BiOCl, SbOCI or a basic salt of 
some other metal, especially titanium and tin. With the exception of the 
titanium precipitate, such basic salts are easily dissolved by filtering and 
treating with 6-normal HCl, or the basic salts of antimony, bismuth and tin 
may be changed into less soluble sulfides by introducing H.S without filtering. 

2. Examine the precipitate of Hg,Cl., PbCl, and AgCl exactly as described 
on p. 283. For the detection of thallium, which is sometimes precipitated 
with this group, consult Part V. 

3. Transfer the filtrate from (1) to a 300 ec. Erlenmeyer flask, fit the latter 
with a two-holed rubber stopper and insert through one hole a right-angled 
glass tube which reaches nearly to the bottom of the flask and through the 
other hole a shorter tube, similarly bent, which reaches only to the bottom 
of the rubber stopper. Raise the longer tube till it is just above the surface 
of the solution, heat the solution to boiling and begin passing a steady stream 
of H.S through the longer tube. -Remove the flame from beneath the flask, 
close the shorter tube with a piece of rubber tubing which has one end stopped 
up with a short piece of stirring rod, and lower the longer tube so that it dips 
well below the surface of the solution. Shake well and continue keeping up 
the pressure of the H:S; in this way the gas is absorbed as fast as the sulfide 
precipitates and the solution is kept saturated with the gas without wasting 
a great deal of it. When the precipitation appears to be complete, shut off the 
HS, open the flask, add an equal volume of cold water and again saturate with 
H.8. Close the short tube and shake the flask well for two or three minutes 
while keeping up the H.S pressure. Finally filter off the precipitate and wash 
it. promptly with H.S water. If an oxidizing agent is present,* considerable 





* If much oxidizing agent or considerable arsenic acid is present, it is best to 
pass SO, into the hot solution until a complete reduction is accomplished and then 


442 SYSTEMATIC ANALYSIS 


free sulfur will be deposited and this will greatly delay the precipitation of the 
sulfides. In case it is desired to know whether a precipitate contains nothing 
but sulfur, wash it several times with alcohol, then with carbon disulfide (away 
from any free flame) and then with alcohol again; this treatment will serve to 
remove the sulfur. It is important to adjust the concentration of the acid 
properly before introducing H.S by measuring the quantity added in getting 
the substance into solution and in precipitating the first group. After diluting 
with water, as above directed, the solution should be about 0.3-normal in acid, 
if more concentrated cadmium and lead will not precipitate and if less con- 
centrated sulfides of zinc, nickel and cobalt may precipitate. On the whole 
it is better to err with too little acid rather than with too much, as enough zine, 
nickel and cobalt will always remain in the filtrate to give a test in the next 


group and the presence of these elements does not seriously interfere with the ~ 


analysis of the second group. Like cadmium, zine gives a white ferrocyanide 
in the confirmatory test for copper and nickel gives a faint blue with ammonia; 
either nickel or cobalt will interfere with the final test for cadmium, but the 
treatment outlined on p. 275 will overcome this difficulty. If after pre- 
cipitating with hydrogen sulfide a turbid filtrate is obtained, due to free sulfur, 
prepare some filter-paper pulp by shaking pieces of filter paper in a bottle 
with hot water, add some of the pulp to the filtrate and filter through a fresh 
filter. The hydrogen sulfide precipitate oxidizes somewhat on being exposed 


to the air, and a little soluble sulfate is likely to form which forms a precipitate — 


on coming in contact with H.S in the filtrate. For this reason the precipitate 
should be washed promptly with hydrogen sulfide water without ever letting 
the filter drain completely until the washing is finished. If the filter clogs, 
place the filter and precipitate in a beaker, shake it up with hydrogen sulfide 
water and filter through a fresh filter. In qualitative analysis, all but the first 
washings of a precipitate should be discarded, as a rule. 

4. Examine the hydrogen sulfide precipitate as directed on p. 273. If 
gold, platinum or considerable tin is present, these elements are often found 
in the residue of mercuric sulfide obtained after treatment with nitric acid. 
When the presence of these elements is suspected, take a little of the residue 
for the mercury test and fuse the remainder in a porcelain crucible with a 
mixture of equal parts potassium cyanide and sodium carbonate. Cool, wash 
out all the soluble alkali salts with water and discard this solution. Gold, 
platinum, tin and lead will be left behind in the metallic condition. Treat 
the metallic residue with dilute nitric acid and test the solution for lead in the 
regular way with sulfuric acid (p. 274). Heat the residue of gold, platinum 


and metastannic acid with concentrated hydrochloric acid. Dilute, filter and 


test for tin with HgCl, in the usual way (p. 276). Dissoive any gold or platinum 
in aqua regia, add ammonium chloride, evaporate to dryness on the water 
bath and treat the residue with a very little water; a yellow precipitate of 
(NH,)2PtCl, shows the presence of platinum. Filter, test with FeSO, for 
gold, and confirm by the charcoal stick reaction (p. 266). 

5. Take a little of the filtrate from (3), boil off the hydrogen sulfide, add 
a little bromine water to oxidize any iron and the last traces of hydrogen 
sulfide and make alkaline with ammonia, If a precipitate forms it may con- 





remove the excess of SO, by a stream of CO.. If the excess of SO, is not removed 
it reacts with H.S as follows: 2H2S+SO.>5 2H,0+38. 


-_ 


eee 


a 


-“j 
; 
“4 
a 
Pe. 
‘ 
a 
. 








ANALYSIS IN THE PRESENCE OF PHOSPHORIC ACID 443 


sist of a phosphate of barium, strontium, calcium or magnesium, or an alkaline 
earth fluoride or oxalate. Phosphoric acid very often occurs in minerals, and 
for this reason a special procedure is often required for the analysis of Groups 
III and IV. Without stopping to filter off the precipitate produced by 
ammonia, dissolve it by the careful addition of a little nitric acid, heat the 
solution nearly to boiling and add an equal volume of ammonium molybdate 
reagent. A yellow precipitate, which may be slow in forming, shows that 
phosphoric acid is present. Arsenic acid gives a similar precipitate (p. 231), 
but arsenic should not be present at this stage in the analysis. If phosphoric 
acid is found present, examine the precipitate as outlined in Table XII. If 
an oxalate or fluoride is indicated by the preliminary examination, especially 
by the test with concentrated sulfuric acid, it is best to remove these acids 
by heating the original substance with concentrated sulfuric acid. This is 
likely to leave, after diluting, an insoluble residue of alkaline earth sulfate 
which should be examined as directed on p. 437. If the behavior of the 
original substance in the closed tube test indicated the presence of non-volatile 
organic matter, it is necessary to remove it before proceeding with the analysis 
of Group III, because tartaric and citric acids, sugars, starches and similar 
substances prevent the precipitation of iron, aluminium and chromium with 
ammonia. Such organic substances can be removed by ignition or by. repeated 
treatment with concentrated sulfuric and nitric acids: adding about 5 cc. of 
sulfuric acid, an equal volume of-concentrated nitric acid and evaporating till 
strong fumes of sulfuric acid are evolved, cooling and repeating the treatment 
with nitric acid as often as necessary. The ignition treatment often makes 
Al,O;, CrsO3, Fe2O; and SiO, insoluble. Fuse such a residue with KHSO, and 
examine it by itself. The treatment with sulfuric and nitric acid is likely 
to leave an insoluble sulfate behind; fuse it with sodium carbonate (p. 109). 

6. If phosphate is found present, analyze as outlined in Table XII, p. 
445. If phosphate is absent, analyze Group III by method A, p. 189 or by 
method B, p. 192. Many of the elements in this group commonly occur in 
different states of oxidation. In reporting the final results of the analysis it 
is not sufficient to state that iron, chromium, or manganese is present, but it 
should be stated in what condition such an element is present in the sample 
as received. It is necessary to determine this by special tests, using the 
characteristic reactions described in Part II under the element in question. 

7. Treat the filtrate from (5) with (NH,),CO; according to Method A, p. 
111 or according to Method B, p. 113. In the former case magnesium will 
not be precipitated with this group and in the latter case it will be precipi- 
tated. In most cases, Method B will be found preferable. 

8. Examine the (NH,),CO; precipitate as directed on p. 111 or as directed 
on p. 113. 

9. Examine the alkali group according to the directions on p. 97 or on 
_ p. 100, omitting the magnesium test provided Method B was used for the 
analysis of Group IV. ‘Test a portion of the original substance for ammonium. 


ANALYSIS IN THE PRESENCE OF PHosPHORIC ACID 


There are three classic methods for examining a solution for the members 
of Groups III and IV when phosphoric acid is present. The first of these 
methods depends upon the fact that when tin is boiled with strong nitric acid, 


4d : SYSTEMATIC ANALYSIS 


insoluble metastannic acid is formed (p. 257) and this precipitate carries down 
with it phosphoric acid, arsenic acid and to a lesser extent other substances — 
such as ferric oxide, titanium oxide, etc. It is probable that an adsorption — 


compound is formed, rather than a simple chemical compound. Instead of ‘- 


using metallic tin, it has been found possible to accomplish the same result “4 
by preparing metastannic acid in advance and adding it to the solution.* 


After the phosphoric acid has been precipitated in this way, it is filtered off s 


and the filtrate examined for Groups III and IV in the usual manner. “7 
The second method of analysis in the presence of phosphoric acid is the 
so-called basic acetate method. If a solution containing ferric iron or aluminium __ 
is carefully neutralized, ferric phosphate or aluminium phosphate is precipitated 
before an insoluble hydroxide or an alkaline earth phosphate is formed. If 
the neutralization goes too far, and an excess of hydroxyl ions is provided, 
ferric and aluminium hydroxides are likely to be formed and alkaline earth 
phosphate will then precipitate. One of the best methods of preventing the 
neutralization going too far is to boil the dilute solution with a little acetic 
acid and considerable sodium acetate. Then, if a slight excess of iron is 
present, all of the phosphoric acid will be precipitated as ferric phosphate and 
the excess of iron will be precipitated as basic acetate. It is better to use fer- — 
ric iron rather than aluminium in this separation, because basic ferric acetate 
by its color shows when enough trivalent metal is present. a 
The third method of analysis is the bariwm carbonate method. This method 
is the same in principle as the basic acetate method except that the solution — 
is neutralized with barium carbonate in the cold. The objection to this method — 
of analysis is that an alkaline earth is added so that a separate portion of the 
solution must be used in the tests for this group. In some special cases calcium 


carbonate, zinc oxide and cadmium carbonate are used instead of barium 


carbonate. All of these substances will neutralize a solution sufficiently to — 


precipitate phosphates and hydroxides of iron, aluminium and chromium. They — 


do not, however, precipitate bivalent metals as a soluble carbonate would.. 


The barium carbonate method is useful for separating the trivalent from — ; 


the bivalent metals of Group III and will be described in Volume I of this 
book. The other two methods will be outlined here. 


PROCEDURE. TIN METHOD 


1. Evaporate the filtrate from Group II just to dryness. Add 10 ce. of + 
concentrated HNO;, evaporate to dryness and repeat the evaporation with 


HNO; once more. Finally add 10 cc. of concentrated HNO; and introduce — 


about 1 gm. of tin foil in small portions. Boil to small volume in order to com- 
plete the precipitation of the tin and pour the concentrated solution into — 
100 cc. of water contained in a narrow cylinder such as a 100-cc. graduate. — 


Next morning siphon off the supernatant liquid and discard the precipitate 


of metastannic acid, which should contain all the phosphoric acid. s 
2. The tin foil usually ‘contains traces of lead and copper. To remove — 
these, saturate the solution with H.S and filter. Reject the precipitate. § 
3. Treat the precipitate with NH,OH and (NH,).S, examine the precipitate _ 
for Group III and the filtrate for Groups IV and V exactly as outlined in — 
Table XI. | 





*Cf. W. MeckLenBura, Z. anal. Chem. 52, 293 (1913) 





BASIC ACETATE METHOD 445 


TABLE XII.—ANAtysis oF Groups III anp IV 1n PRESENCE or PHOSPHATE 














Tin Method. Basic Acetate Method. 

Remove HS and HCl from filtrate from | Treat filtrate from Group II with NH,OH 
Group II, and evaporate repeatedly with | and (NH4)S. Filter and keep filtrate. 
HNO;. Add tin foil, concentrate, dilute | Examine the precipitate according to Table 
and allow the precipitate to settle. (1) VI up to and including treatment with 

. HNO; and KCI0O;3. Test for Fe+++ in 
filtrate from MnO, precipitate. Add 
NH.OH, HC.H;0z, FeCls and NH.C2H:02. 
Dilute and boil. (4) 

Precipi- | Solution: Groups III, IV | Precipitate: | Filtrate: Nit+t, Co,++ 
tate: and V. Saturate with| FePO, and (Zn*+)Batt, §rtt, 

(H2SnO;)x-| HS and filter. (2) Fe(OH).C2H;02.| Cott, Mgtt. Add 
(P20s)y. . Reject. NH,OH- and saturate 

Reject. with HS. (5) 

Precipi-| Solution: Precipi- | Filtrate: 
tate: CuS, | Groups III, CoS,NiS, | Batt, 
PbS from | IV and V. (ZnS). Srtt, 
impurities | Add Examine | Catt, 
in tin foil.| NH,zOH as in Ta-| Mgtt. 
Reject. and . ble VI. Add this to 

(NH4)28S the filtrate 

and con- obtained 

| tinue as in after the 

Table XI. jirst treat- 

(3) ment with 
NH,OH 
and 
(NH,4)8S 
and exam- 
ine for 
Groups IV 
and V as 
in Table 
XI. (6) 




















Basic AcETATE METHOD 


4, Treat the filtrate from Group II with NH,OH and (NH,).S in the usual 
way. The precipitate may contain FeS, NiS, CoS, ZnS, MnS, Al(OH) , AlPOx,, 
Cr(OH)s, CrPOQ,, Bas(PO,)s, Srs(PO,)s, Ca;(POx)2, and MgNH,POs,. The 
filtrate may contain Bat*, Srt*+, Cat+, Mgt+, Nat, Kt and NH,*. Set 
the filtrate aside to be combined later with a solution which will contain 
any Batt, Srt*, Cat* and Mg++ that was precipitated as phosphate by 
NH.OH. 

Dissolve the precipitate in HCl, adding a little HNO; if necessary. Dilute, 
filter off any residual sulfur and evaporate the filtrate nearly to dryness to 
remove the excess acid. Dilute to 20 cc., make alkaline with NaOH and add 
more water if a very bulky precipitate is formed. Add about 2 cc. of Na,O0: 


446 - SYSTEMATIC ANALYSIS 


powder in small portions while stirring the cold solution. Then add 5 ce. of 
4-normal Na2CQ; solution and boil to decompose the excess of NazOz. Cool, 
dilute with an equal volume of water and filter. The filtrate will contain all 
the aluminium as sodium aluminate, all the chromium as sodium chromate, 
usually most of the zine as sodium zincate and some or all of the phosphoric 
acid as sodium phosphate. Examine this filtrate for aluminium, chromium and 
zinc exactly as outlined in Table VI, p. 192. The presence of the phosphate does 
no harm, as zinc phosphate is readily soluble in ammoniacal solutions. When 
the solution which was treated with sodium hydroxide and sodium peroxide 
contains alkaline earths, the phosphates of these elements are precipitated by 
the alkali and if the solution contains no alkaline earth metal, usually most of 
the phosphate remains in solution. The sodium carbonate is added with the 
sodium hydroxide and sodium peroxide to ensure the complete precipitation 
of alkaline earth and to prevent the precipitation of an alkaline earth chromate. 

Dissolve the precipitate produced by NaOH, Na,O. and NazCO; in 15 ce. 
of 6-normal HNO;, adding HO. a few drops at a time till precipitate is all 
dissolved. Filter if any filter paper is present and evaporate the filtrate nearly 
to dryness. Add 16-normal HNO; and treat with KCIO; exactly as described 
on p. 193. If a precipitate of MnO, is formed, filter it off and, if deemed 
necessary, confirm the manganese in the usual way. Take one-tenth of the 
filtrate from the KCIO; treatment, evaporate it just to dryness, moisten the 
residue with a few drops of concentrated HCl, dilute with 5 cc. of water and 
test for iron with KCNS. Neutralize the remainder of the solution with 
ammonia and continue adding ammonia until a permanent precipitate is formed 
or the solution becomes alkaline. Dissolve the precipitate in acetic acid, 
avoiding an excess, and add 15 cc. of 3-normal NH,C,H;0, solution. If the 
solution is then of a rich brownish red color, it is evident that more than 
sufficient iron is present to combine with all the phosphoric acid. Otherwise 
add FeCl; solution, drop by drop until such a color is produced. Dilute to 
at least 100 cc. and hoil for five minutes in a 500-cc. flask. If a large precipitate 
is formed, dilute to 250 cc. with hot water and boil a minute longer. Allow 
the precipitate to settle for a minute or two, filter while still hot through a 
plaited filter and wash the precipitate with hot water containing a little 
ammonium acetate. Add 10 cc. more of NH,C.H;0; to the filtrate, again boil 
and collect any further precipitate on a separate filter. Reject both of these 
precipitates, which contain all the phosphoric acid as pale yellow ferric phos- 
phate and the excess of the iron as basic ferric acetate. 

5. Make the filtrate ammoniacal and saturate it with H.S. Filter and 
examine the precipitate for nickel, cobalt and zinc according to the directions 
on p. 193. 

6. Add this last filtrate, which may contain Batt, Sr++, Cat* and Mgtt 
as well as the added NH, salt, to the filtrate obtained by the original treat- 
ment with NH,OH and (NH,).S. Examine the combined filtrates for Groups 
IV and V as outlined in Table XI, p. 440. 


EXAMINATION FOR THE NEGATIVE ELEMENTS (ANIONS) 


The tests for the acids (anions) are usually made after the analy- 
sis for the metals (cations); the preliminary examination (heating in 
the closed tube and with dilute and concentrated sulfuric acid) and 


_ EXAMINATION FOR THE NEGATIVE ELEMENTS 447 


the solubility, combined with the knowledge of the metals that are 
present, suffice to tell us what acids may and what acids may not be 
present. 

In order to avoid side-reactions, the acids are usually obtained in 
the form of the neutral alkali salts before proceeding to test for 
them. 


PREPARATION OF THE SOLUTION FOR THE ANALYSIS FOR ACIDS 


Two cases may be distinguished: 


A. The original substance contains no heavy metal (i.e., only alkalies 
or alkaline earths are present). 

(a) The substance is soluble in water. 

Test the solution with litmus paper to see whether it is acid, alkaline, 
or neutral. 

An Alkaline Reaction shows the possible presence of alkali cyanides, alkali 
nitrites,* borates, tertiary phosphates, alkali sulfides, thio salts of the alkalies, 
alkali silicates, etc. 

An Acid Reaction is shown by many acid salts (cf. p. 49). 

Divide the solution into two parts. If it is neutral, analyze it directly for 
the acids; if it is alkaline, neutralize + half of it with acetic acid and the 
other half with nitric acid; if it is acid, neutralize with sodium carbonate 
solution. 

(b) The substance is insoluble or very difficultly soluble in water, but 
readily soluble in dilute acids. In this case only the acids of Groups III and 
IV need be tested for. 

Boil the dry substance with a little concentrated sodium carbonate solution 
and filter. The filtrate contains the acids in the form of their sodium salts. 

Neutralize the solution with dilute nitric acid. 

(c) The substance is insoluble in water and in dilute acids. 

The following substances may be present: BaSQ., SrSOu, (CaSO,), CaF», 
and silicates, which often contain salts of H;PO., HBO:, H.SO,, HF, and HCl. 

Fuse the substance with sodium carbonate in a platinum crucible, extract 
with water, and use the aqueous solution thus obtained for the analysis for 
acids, after neutralizing. 

If the substance is partly soluble in water and in acids, first treat with water 
and then with sodium carbonate solution and fuse the residue with solid 
sodium carbonate. Analyze separately the three solutions thus obtained. — 

B. The substance contains heavy metals. 

(a) It is Soluble in Water or in dilute acids, and contains no non-volatile 
organic matter (no carbonization in the closed-tube test.) 

Treat the solid substance with suificient concentrated sodium carbonate 





* Perfectly pure alkali nitrites are neutral. The alkaline reaction of the com- 
mercial salts is due to the presence of alkali oxide or silicate. 

{ Taiosalts, silicates, stannites, stannates, aluminates, molybdates, tung- 
states etc., will yield precipitates at this point which should be examined accord- 
ing to Table XIII. 


7 


448 -  §YSTEMATIC ANALYSIS - 


solution to make the resulting solution weakly alkaline, and filter. If ammo- — 
nium salts are present, first boil it with the solution of sodium carbonate until — 
the vapors from the solution no longer smell of ammonia, and then filter. “all 

Divide the resulting solution into two parts, making one part acid with — 
acetic acid, and the other with nitric acid. a 

(b) The Substance is Soluble in Water or Dilute Acids and Contaiuas | 
Non-volatile Organic Matter.—If the metals of the ammonium sulfide and 
hydrogen sulfide groups are both present, pass hydrogen sulfide into the 
weakly acid solution until it is saturated, filter off the precipitate, add ammonia 
to the filtrate until it is slightly alkaline, filter again, and make this last filtrate — 
acid with acetic acid and evaporate to a small volume. Filter off the deposited 
sulfur, treat the solution with solid potassium carbonate, filter if necessary, 
carefully acidify with nitric acid, stir vigorously, and if any potassium-acid — 
tartrate is formed, filter it off and test.as described on p. 363. Test the — 
filtrate for the remaining acids. a 

(c) The Substance is Insoluble in Strong Acids.—Besides the salts’ men- 
tioned under A (c), the following may be present: AgCl, AgBr, AgIl, AgCN, 
PbSO,, silicates (ferro- and ferricyanides). 

If silver is present, the halogen acids must be looked for. Reduce the = 
insoluble silver salt by zinc and sulfuric acid, filter off the residue, and examine - 
the filtrate according to Table XV for HCl, HI, HBr, and HCN. ; 4 

If the insoluble substance contains lead, boil it with sodium carbonate — 
solution and filter; make the filtrate acid with hydrochloric acid, and test with _ 
BaCl. for H.SQ,. : 

If silicates are present, H;PO., HF, HBO., HCl, and H.SO, must also be — 
tested for. _ 

In whatever way a solution is prepared, determine its behavior toward | 
silver nitrate and barium chloride in order to ascertain to what groups the ¥ 
acids present belong. 3 

Then make the necessary tests for the individual members. 





_ 


EXAMINATION FOR THE NEGATIVE ELEMENTS 


449 


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ISH], AAIUOTHD WAMUVG THI — ATX AVI, 


450 


EXAMINATION FOR THE NEGATIVE ELEMENTS 451 


TABLE XV.—EXAMINATION OF Group I 


First test for HCN by placing a little of the solution on a watch-glass, adding a few 
drops of yellow ammonium sulfide, evaporating carefully to dryness, acidifying 
the dry mass with HCl, and adding a drop of FeCls solution. If a blood-red col- 
oration is produced, HCN is present, in which case treat a larger portion of the 
neutral solution with nickel sulfate * solution in excess, and filter. 





PRECIPITATE. SOLUTION. 
Ni(CN)2 Treat the solution, which is now free from hydrocyanic acid, with a 
Discard. a little caustic soda solution (free from halogen), boil and filter off 


the precipitate of Ni(OH)>2, divide the filtrate into two parts and use 
one part for the HBr and HI tests and the other for the HCl test. 





Tests for HI and HBr. 


Test for HCl. 








Make the solution acid with dilute 


H.SO;, add chlorine water drop 
by drop, and shake the solution 
with CS: or CHCl;. If the latter 
is colored violet, HI is present. 
By further addition of chlorine 
water, the CS, or CHCl; ts de- 
colorized completely if HBr is 
absent, but turned  yellowish- 
brown if HBr is present. If too 
much chlorine water is used a 
wine-yellow color is produced. 





Make the solution slightly acid 
with HNO;, and add dilute 
AgNO; drop by drop. AglI 
and AgBr are first precipitated 
(yellow). Filter and add more 
AgNO;. If the precipitate 
still appears yellow, filter 
through a new filter, and 
again add AgNO; to the filtrate 
until a white precipitate of 
AgCl is formed if HCl is 
present. 





* If ferricyanic acid is present also add a little ferrous sulfate. 


precipitated by nickel sulfate. 


Group II 


Ferrocyanic acid is completely 


The members of this group are almost always detected in the preliminary 
examination. The special tests for these acids are described on p. 329 et seq. 


Group III 


SO2, COz, H2C.0, are recognized in the preliminary examination. HPOs, 
H,P,0,, HBO,, and H.C,H,O, are tested for separately by the special reactions 
described on p. 347 et seq. 


Group IV 


Cr0;, H;PO,, and H.8.0; are detected in the preliminary examination, and 
in the analysis for metals. 


Group V 


HClO; and HNO; are usually detected in the preliminary examination. 
Their presence is, however, always confirmed by the procedure described on 


p. 398. 


Groups VI and VII 


These acids are usually detected in the preliminary examination. Their 
presence is confirmed by the tests described under H.SO,, HF, and SiOx. 


452 SYSTEMATIC ANALYSIS 


B. THE SUBSTANCE IS A METAL OR AN ALLOY 


The analysis of a metallic alloy is much simpler than that of a mixture of 
salts, because there are no acids to test for. Of the electro-negative elements, 
usually only phosphorus, silicon, carbon, and sulfur have to be considered. 

As all metals, with the exception of gold, platinum, tin, and antimony, are 
soluble in nitric acid, alloys are usually brought into solution by dissolving 
therein, and the use of aqua regia is only necessary in a few cases. Many 


alloys rich in silicon (e.g., copper silicide) are extremely difficultly soluble even 4 


in aqua regia, and are best brought into solution by fusing with caustic alkali. 


in a silver crucible, and afterwards dissolving the melt in nitric acid. 

It is not advisable to dissolve an alloy in hydrochloric acid, for phosphides, 
carbides, silicides, sulfides, and arsenides, which are often present in small 
amounts, are decomposed by this acid in such a way that the negative elements 
are evolved as hydrogen compounds, and thus escape detection. For the 
analysis of ordinary alloys, the following procedure is used: 

Place 1 or 2 gms. of the alloy (best in the form of borings) in a 200-ce. porce- 
lain dish and treat them under a good hood ‘with about 20 ce. of nitric acid, sp. gr. 
1.25-1.30 (1 vol. cone. HNO;+1 vol. H:O). After the first violent reaction 


is over, carefully evaporate (with constant stirring) almost to dryness, being — 


careful to avoid overheating; * add a little water and heat. 

(a) The mass dissolves completely. The alloy contains neither tin nor 
antimony; analyze it according to Table XVI. 

(b) The mass does not dissolve completely, but a white, greenish residue 
remains; analyze according to Table XVII. 


C. THE SUBSTANCE IS A LIQUID 


The color, odor, and reaction towards litmus enable one to draw important 
conclusions. 


(a) The solution reacts neutral; it contains no free acid, free base, acid 


salt, no salt which shows an acid or alkaline reaction on account of hydrolysis, 
nor any insoluble salt. 


First of all, determine whether there are any solid substances dissolved in — 


the liquid by evaporating a small portion to dryness at as low a temperature 
as possible (so as not to lose any volatile substances). If a residue remains, 
examine it according to A, p. 428. 


(b) The solution reacts alkaline. An alkaline reaction may be due to the 


presence of hydroxides of the alkalies or alkaline earths, peroxides, carbonates, 
borates, cyanides, silicates, sulfides (zincates, aluminates, molybdates, tung- 
states) of the alkalies, as well as by ammonia or hypochlorites, ete. 

If the solution, for example, contains hydroxides or carbonates of the alkalies, 
it is evident. that substances which are precipitated by them cannot be present 
at the same time, except, in some cases, in the form of complex ions (cyanides, 
tartrates, etc.) 

At once test the solution for peroxides, hydroxides, and carbonates, as well 
as for the sulfides of the alkalies. 





* Otherwise insoluble basic salts are likely to be formed. If this be the case, — 


as is often shown by the dark color of the residue, add a little conc. HNOs, heat 
the liquid somewhat and then dilute with water. 





453 


THE SUBSTANCE IS A METAL OR AN ALLOY 


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SYSTEMATIC ANALYSIS 


454 


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THE SUBSTANCE IS A METAL OR AN ALLOY 455 


To test for peroxides * (H.0.) heat a little of the solution with a few drops 
of cobalt nitrate solution; a black precipitate shows the presence of H2O:.t 
Or test the solution by adding some titanium sulfate solution and acidifying 
carefully with cold dilute sulfuric acid; a yellow coloration shows the presence 
presence of H,0..t 

A still more sensitive reagent, according to Schéne,§ is a very dilute solu- 
tion of FeCl;+K;[Fe(CN),]. If the slightest trace of H.O,. is present in the 
solution, the red solution becomes greenish, and after a time Prussian blue 
separates out. 

In order to detect the presence of hydroxides and carbonates in the presence 
of H;02, boil a portion of the solution for a long time in a porcelain dish in order 
- to destroy the peroxide, and then add barium chloride until no more precipitate is 
formed. If the solution now shows an alkaline reaction, the presence of hydrox- 
ides || is assured. If the precipitate produced by barium chloride dissolves 
in acid with effervescence, and the gas evolved renders barium hydroxide 
solution turbid, carbonates are present. If the solution smells of ammonia, 
evaporate a small portion to dryness in order to see whether other compounds 
are present, and examine the residue according to A, p. 423. 

(c) The solution reacts acid; it can then contain substances which are 
soluble in water and in acids, as well as free acids. Evaporate a small portion 
to dryness in order to see whether any non-volatile matter is present. If no 
residue is obtained, neutralize the solution with soda and test for acids. If 
a residue is obtained, examine it according to A, p. 423. 

D. The substance to be analyzed is a gas. 

This case will be considered in Volume II under Gas Analysis. 





* See foot-note, p. 428. 

{If the alkaline solution contains hydrochlorites or sulfides, these will also 
give a black precipitate with cobalt nitrate; the above reaction serves only to detect 
HO, in the absence of hypochlorites or sulfides. If the latter bodies are present, 
H,O2 cannot be present, because hypochlorites are reduced to chlorides and sul- 
fides oxidized to sulfates by H2O.. f 

The presence of hypochlorites is usually detected by. the odor, on acidifying 
with dilute H.SO,, the odor of chlorine can be detected. Sulfides give off H.S on 
being acidified. Hypochlorites and su'fides cannot exist together in the same 
solution. . 

t H.O2 can also be detected by the chromic acid reaction, but this test is less 
certain than that with titanium sulfate. | 

§ Ber., 7, 1695. 

|| Either as such in the original solution or by hydrolysis of peroxides. This 
test for OH ions in the presence of carbonates has been used for years by the 
author. It is reliable, though care must be taken to add an excess of BaCh, as 
otherwise an alkaline reaction may be due to BaCO;, which is more soluble in 
water than in BaCl, solution. 


PART V.—REACTIONS OF SOME OF THE RARER 
METALS 


In the treatment of the rare metals, the same method and order 
will be used as with common metals. 
THE ALKALI GROUP 
CASIUM, RUBIDIUM, LITHIUM 
CESIUM, Cs. At. Wt. 132.81. M. Pt.=26° 


Occurrence.—Cesium and rubidium are not rare elements, strictly. 


speaking, for they are found almost everywhere, but always only 


in very small amounts. Thus cesium replaces potassium in many 
feldspars and micas, and is found in many rocks which carry these 
minerals, as well as in mineral waters which ooze from them. Czsium _ 
and rubidium were discovered in the mother liquor of Durkheimer 
brine in the year 1860 by Bunsen and Kirchhoff by means of the spec 
troscope. ae 

Pollucite, a mineral closely related to leucite, found at Elba and ~ 
crystallizing in the regular system, is a typical cesium mineral. Its 
composition is H2@s4Al4(Si03)9. 

Cesium and rubidium in all their reactions behave almost exactly 
like potassium. The principal difference between these three metals 


lies in the different solubilities of their corresponding salts, as will @ 


be seen by the table given on p. 459. 


REACTIONS IN THE WET WAY 


A solution of cesium chloride should be used. } 

1. H2[PtCls] produces a yellow, crystalline precipitate which is — 
of a lighter color than the corresponding potassium salt and is much — 
less soluble; 100 parts of water dissolve at 0° C. only 0.024 part and at 
100° C. 0.377 part of the salt. 


2. Tartaric Acid produces, as with potassium and rubidium, a a 


white, crystalline precipitate, CsHC4H40.; 100 parts of water dis- — | 
solve at 25° C. 9.7 parts, and at 100° 97.1 parts of the salt. 9 


3. He[SnCl¢] (a solution of SnCl, in concentrated HCl) produces — 


in concentrated solutions a white, crystalline precipitate of Cs2{SnCle] q 
456 





THE ALKALI GROUP 457 


(octahedrons). Ammonium salts give the same reaction, but potas- 
sium and rubidium salts do not. 


REACTIONS IN THE DRY WAY 


Cesium compounds color the flame-reddish violet, very similar 
to the potassium flame. 

Flame Spectrum. An intensely blue double line, 455.5upu, 459.3up. 
At higher temperatures a number of paler lines appear of subordinate 
importance; in the red 697.3uu and 672.3uy, in the orange yellow 
621.3uu and 601yup, in the yellow 584.5yy, in the green 566.4uu, 563.5up, 
550.3uu, 547.1uy, 541.9un, and 535.luu. Moreover, a faint. contin- 
uous spectrum is seen from the yellow to the blue. See chart. 


RUBIDIUM, Rb. At. Wt. 85.45. M. Pt.=38° 


Occurrence.—Rubidium almost always accompanies cesium and is 
found in many mineral waters; in carnallite from Stassfurt; in lepido- 
lite, (Li,K,Na)2Ale(F,OH)2Sis09; in triphylite, (FeMn)(LiCsRb)PO,; 
and in spodumene, (LiNa)AIl(SiO3)2 a mineral of the pyroxene group. 
Lepidolite from Rozena contains about 0.54 per cent Rb and 0.0014 
per cent Cs. As far as the author knows, there is no typical rubidium 
mineral. 


REACTIONS IN THE WET WAY 


1. H»[PtCls] produces, as with cesium and potassium salts, a 
white, crystalline precipitate of Rbe[PtCle], which is more difficultly 
soluble than the corresponding potassium salt, but more soluble 
than the cesium salt; 100 parts of water dissolve at 0° C. 0.134 part, 
and at 100° 0.634 part of the salt. 

2. H2[SnCle] produces a. white precipitate only in very concen- 
trated solutions. The salt is much more soluble than the corre- 
sponding cesium salt, but this reaction is not suitable for separating 
the two metals. 

3. Tartaric Acid produces a precipitate of RbHC.zH40¢ only in 
concentrated solutions; 100 parts of water dissolve at 25° C. 1.18 
parts, and at 100° C. 94.1 parts of the salt. The corresponding cesium 
salt is more soluble, while that of potassium is less soluble. 


REACTIONS IN THE DRY WAY 


Flame Coloration.—Similar to csesium. 
Flame Spectrwm.—Violet double lines 420.2uu and 421. Bias also 
the red double line 781.1uu and 795.0uy. At higher temperatures a 


458 REACTIONS OF SOME OF THE RARER METALS 


continuous spectrum is visible from the yellow to the blue in which 
the following lines are to be found in addition to those given above; 
in the orange-yellow 62).8uu, 626.1uu, 620.64, 617.7up; in the yellow- 
green 572.4uu, 570.0uu, 564.8un and of subordinate importance the 
ereen lines 543.5upu, 536.5uu, and 516.8up. 


LITHIUM, Li. At. Wt. 6.94. M. Pt.=186° 


Occurrence.—Lithium is found to a greater extent in nature than 
cesium and rubidium; in triphylite, (Fe,Mn)(Li,Cs,Rb)PO4; in 
petalite, Al(Li,Na,H)Si40i0, a mineral of the feldspar group (also 
called castorite); in amblygonite, Li(AIF)PO4, monoclinic; in lepido- 
lite, Ale(Li,K,Na)o(F,OH)2Siz09; in many varieties of tourmaline, 
and muscovite, in epidote and orthoclase, and consequently in many 


mineral-spring waters. In some cases as much as 36 mgs. Li are 


contained in a liter of spring water. 

Lithium is the lightest of all metals, and floats on petroleum. 
It oxidizes quickly in the air, and decomposes water at ordinary tem- 
peratures, forming LiOH, which dissolves slowly in the water; the 
solution reacts alkaline and absorbs carbon dioxide from the air 
with avidity, forming difficultly soluble LigCOs3. 


Lithium chloride is soluble, even in the anhydrous state, in a mix-— 


ture of alcohol and ether as well as in amyl alcohol (difference from 
the remaining metals of this group). 


REACTIONS IN THE WET WAY 


A solution of lithium chloride should be used. 

1. He{PtCle] produces no precipitation. 

2. Tartaric Acid produces no precipitation. 

3. NazHPO, produces from boiling, moderately concentrated solu- 
tions a white precipitate of tertiary lithium phosphate. ‘lhe pre- 
cipitation is only quantitative when the solution is made alkaline 
with caustic soda, evaporated to dryness, and taken up in water con- 
taining ammonia: 


HPO4+3Lit+OH~ — LisPO4+ H20. 


Lithium phosphate is fusible (difference from magnesium and 
the alkaline earths). . 
4. (NH4)2CO3. If ammonium carbonate and ammonia are added 


to. a concentrated lithium solution, lithium carbonate is precipitated 


in the form of a white powder. The salt, contrary to the other alkali 





THE ALKALI GROUP 459 
carbonates, is: very difficultly soiuble in water: 100 parts of water 
dissolve at 13° C. 1.31 parts of LigCO3. In the presence of con- 
siderable alkali chloride or of ammonium chloride, no precipitation 
' takes place. 


REACTIONS IN THE DRY WAY 


Flame Coloration.—Pure lithium salts impart a magnificent, car- 
mine-red coloration to the gas-flame. If considerable amounts of 
sodium salts are present at the same time, the lithium flame is com- 
pletely masked; but if the flame is observed through cobalt glass the 
red color becomes distinctly visible. 

Flame Spectrum.—An intensely red line 670.8up and at higher 
temperatures the pale orange-yellow line 610.8uy. 


SUMMARY OF THE ALKALI METALS 








LITHIUM. Sopium. Porassium.} Ruprprium.| Casrium. 
Atomic WEIGHT..... 6.94 23.00 39.10 85.45 132.81 
MELTING-POINT....... 180° 95.6° 62.5° 38 .5° 26-27° 
X.[PtCl,]: 
Solubility in alcohol soluble soluble insoluble | insoluble | insoluble 
odes sp aes considerable | considerable} 1.12 0.141 0. 079 
ah 100s. s considerable | considerable | 5.13 0.634 0.377 
XHC,H.Os: 
Pe eae 1 i sat considerable | considerable | 0.425 — a 
1 he? lal ea _ — — 1.18 9.7 
ALUMS: 
aE er bul a \ ne = 13.5 2.27 0.619 
CHLORIDES: ‘ 
Sapte = oa soluble insoluble | insoluble | insoluble | insoluble 
CARBONATES: : 
age eee ome insoluble insoluble | insoluble | insoluble | soluble 




















DETECTION OF LITHIUM, RUBIDIUM, AND CASIUM 
IN THE PRESENCE OF CONSIDERABLE SopIuM AND POoTAssIUM 


Evaporate the solution containing the metals as chlorides almost to dryness, 
triturate the residue with 90 per cent alcohol and filter. The alcoholic solution 
contains all of the Li, Rb, and Cs, with small amounts of K and Na. Evaporate 
again nearly to dryness and once more extract with alcohol and filter. If only 
traces of the rare alkalies are present, the residue must be extracted several 
times with alcohol. 

Evaporate the alcoholic extract to dryness and treat the residue with con- 


/ 


460 REACTIONS OF SOME OF THE RARER METALS 


centrated hydrochloric acid. This is done in order to change LiOH into LiCl. — 
Some of the former is formed by evaporating the solution, and it is insoluble — 
in alcohol-ether. Evaporate once more, gently ignite the residue over the free. 
flame, and, after cooling, triturate the residue with ether-aleohol mixture, using 
a elass rod, and filter through a filter that is wet with the ether-alcohol mixtures . 
The filtrate contains the lithium. Evaporate it to dryness and test for lithium | 4 
by means of the flame reaction or in the spectroscope. a 

Dissolve the residue insoluble in alcohol-ether in a little water, treat with — 
chloroplatinic acid and filter. Extract the precipitate repeatedly with small — 
portions of boiling water, decanting off the liquid each time through the filter. | 

The potassium chloroplatinate dissolves in the hot water, forming a — 
yellow solution. Continue the treatment with hot water until the residue is — 
of a light-yellow color. Dry, place in a porcelain boat, and heat in a glass — 
tube, made of difficultly fusible glass, in a stream of dry hydrogen; in this — 
way the alkali chloroplatinates are reduced to chloride and platinum: 


X{[PtCh]-+2H2 =4HCl ¢ +2XC1+Pt. 


After cooling, treat the residue with a little water, filter off the platinum, — 
evaporate the solution to dryness and test in the spectroscope for: Cs and Rb. | 

In order to detect lithium, cesium, and rubidium in a silicate undecom-— 
posable by acids (lepidolite, for example) decompose the finely powdered — 
mineral with hydrofluoric and sulfuric acids as described on p. 419, change 
the sulfates to chlorides by the addition of barium chloride, and ‘free the > 
solution from other metals as described on p. 96, and then carry out the 
above separation. | 





METALS OF THE (NH,),S GROUP 


BERYLLIUM, ZIRCONIUM, THORIUM, YTTRIUM, ERBIUM, CERIUM, 
LANTHANUM, DIDYMIUM, TANTALUM, NIOBIUM 


BERYLLIUM, Be. At. Wt. 9.1. M. Pt.=>1880° 


Occurrence.—Chrysoberyl, Be(AlOz)2; phenacite, Be2SiO4; _bery], 
BezAleSigO1g; euclase, AlBeHSiOs; meliphanite, Be2CazNaSizO10F; 
and leucophanite, BeCaNaSi2OcF. 

Beryllium is a bivalent metal, and forms a white oxide, BeO, 
which is soluble in acids. Beryllium salts react acid in aqueous solu- 
tion and possess a sweetish, astringent taste. The element is often 
called Glucinum, Gl, in England and the United States. 


REACTIONS IN THE WET WAY 


Use a solution of BeSO4-4H20. 

1. Ammonia and Ammonium Sulfide produce a white precipitate 
of Be(OH)2, similar in appearance to AlfOH)s, insoluble in an excess 
of the precipitant, but readily soluble in HCl, forming a colorless 
solution. The yellow color often obtained in dissolving the hydroxide 
in hydrochloric acid is due to traces of ferric chloride. 

2. KOH precipitates white, gelatinous beryllium hydroxide, readily 
soluble in an excess of the reagent, forming potassium beryllate: 


Be**++20H- — Be(OH)2; Be(OH)2+20H~ — [BeO2]~+2H20. 


The alkali beryllates are decomposed hydrolytically on boiling their dilute 
aqueous solutions, all of the beryllium being precipitated as hydroxide. The 
precipitate thus obtained is denser than that thrown down by ammonia, 
and differs from the latter by being insoluble in potassium carbonate and 
difficultly soluble in ammonium carbonate solutions; it is also much more 
difficultly soluble in dilute acids. A solution containing a considerable excess 
of alkali hydroxide does not give a precipitate of beryllium hydroxide by 
boiling. 

3. Ammonium Carbonate produces a white precipitate of beryl- 
lium carbonate, readily soluble in an excess of the reagent (difference 
from aluminium); by boiling the solution, the beryllium is precipi- 
tated as white basic carbonate. This property enables us to sepa- 
rate beryllium from iron and aluminium. The separation is not 

461 


462 REACTIONS OF SOME OF THE RARER METALS 


sharp, however; which can also be said of the separation by means of 
caustic potash. 


In order to make a quantitative separation, the beryllium hydroxide or 7g 


carbonate must be redissolved and the precipitation repeated several times. 


4. BaCO3 precipitates beryllium completely in the cold as hydroxide. — 


5. Oxalic Acid and Ammonium Oxalate cause no precipitation 
(difference from thorium, zirconium, erbium, yttrium, cerium, lan- 
thanum, and didymium). 

6. KeSO4 gives with beryllium sulfate a beautifully crystalline 
double salt, Ke2Be(SO4)2-2H20, which is soluble in a concentrated 
solution of KeSOx4 (difference from Ce, La, and Di). | 

7. BeCle is soluble in a mixture of equal volumes of saturated 
aqueous and ethereal hydrochloric acid, while the hydrous alumin- 
ium chloride is not (good method for separating Be and Al).* 

There are no characteristic dry reactions for beryllium. 


ZIRCONIUM, Zr. At. Wt. 90.6. M.Pt.<Si 


Occurrence.—Zircon, Zr28iO4, tetragonal, isomorphous with rutile 
TizO4, thorite (orangite), ThSO., cassiterite, Sn2O4, polianite, 
Mn204, and plattnerite, Pb2O4; baddeleyite, ZrO2, monoclinic. 

Zirconium forms two oxides: Zirconium dioxide, ZrOez, and azir- 
conium pentoxide, Zr2O5. The former is the more important and can 
be dissolved by heating for a long time with a mixture of two parts 
of concentrated H2SO4 and one part of water and afterwards diluting. 


The mineral zircon, ZrSiO., cannot be decomposed by such treatment. It 
must be finely pulverized and fused with four times as much sodium carbonate 
at a high heat in a platinum crucible; sodium silicate, Na,SiO., and sodium 
zirconate, NasZrQO,, are formed. On treating the melt with water, the former 
salt dissolves, while the latter is decomposed hydrolytically, forming sodium 
hydroxide and sandy, insoluble zirconium hydroxide; the latter retains some 
of the caustic soda very persistently. After washing the residue, heat it, without 
previous drying, with concentrated sulfuric acid at a temperature near 
boiling-point; in this way anhydrous Zr(SO,)2 is obtained. By pouring 
water over the latter, the salt Zr(SO,)2-4H2O0 is formed, which dissolves slowly 
in cold water, but more readily in hot water, forming a solution with an acid 
reaction. 


REACTIONS IN THE WET WAY 
A solution of zirconium nitrate, or a freshly prepared one of are 
oxychloride, may be used for the following reactions: 


1. NH4OH and (NH4)2S produce a white gelatinous precipitate 
of Zr(OH)a4, insoluble in an excess of reagent. 





*F.S. Havens, Z. anorg. Chem., 18 (1898), 147. 


; 
‘ 
! 
: 
| 





METALS OF THE (NH,).S GROUP 463 


2. KOH and NaOH likewise produce the same precipitate insol- 
uble in an excess of reagent (difference from Al and Be). When the 
zirconium hydroxide is produced in the cold it is readily soluble in 
dilute acids; but when thrown down from a boiling solution it is very 
difficultly soluble in dilute acids, though it will dissolve even then in 
concentrated acids without difficulty. 

3. (NH1)2CO3 produces a white, flocculent precipitate of basic 
carbonate, readily soluble in an excess of the reagent, but reprecipi- 
tated by boiling. 

4. K2COs and NazCOz produce white precipitates somewhat 
soluble in an excess, but reprecipitated by ammonia. 

5. BaCO3 causes incomplete precipitation, even on boiling. 

6. Oxalic Acid gives a white, ftocculent precipitate of zirconium 
oxalate, readily soluble in an excess of oxalic acid, difficultly soluble 
in dilute hydrochloric acid, and readily soluble in ammonium oxa- 
late. From the solution in (NH4)2C2O4 the zirconium is not pre- 
cipitated by the addition of dilute HCl (difference from Th). 

7. Ammonium Oxalate behaves the same as oxalic acid. 

From the solution in ammonium oxalate, zirconium is not pre- 
cipitated on the addition of hydrochloric acid (difference from tho- 
rium). 

Remark.—A solution of zirconium sulfate behaves quite differently from 
that of the nitrate and oxychloride towards oxalic acid and ammonium oxalate, 
a fact which, although published by Berzelius and also by Pfaff, had been 
entirely forgotten by most chemists until their attention was called to it by 
R. Ruer.* 

On treating an aqueous solution of zirconium sulfate with oxalic acid or 
ammonium oxalate there is no precipitation; in fact, precipitation will not 
take place from nitrate or chloride solutions when these contain sufficient 
sulfuric acid, sodium or ammonium sulfate. 

The cause of this different behavior lies in the fact that zirconium forms 
complex compounds with sulfuric acid and alkali sulfates. Thus the solution 
of zirconium sulfate contains the acid H,[ZrO(SO,).], and on treating a solution 
of the oxychloride or nitrate with sodium or ammonium sulfate (but not the 
potassium salt) the sodium or ammonium salt of this complex acid is formed: 


ZrOCl,+2Na80, =2NaCl+ Na,[ZrO(SO,),]. 


These compounds, however, are electrolytically dissociated in aqueous 
solution as follows: 


H,[ZrO(SO,)] = 2H*++ [ZrO(SO,).]*. 
As the zirconium is present in the anion it ‘erint react with oxalic acid. 
8. HF causes no precipitation (difference from Th and Y). 
* Z. anorg. Chem., 42, 85 (1904). 





464 REACTIONS OF SOME OF THE RARER METALS 


9. K2SO4.--A concentrated, cold solution of K2SO4 precipitates — 
little by little all of the zirconium as potassium zirconium sulfate, — 
insoluble in an excess of the reagent (difference from Al and Be). — 
The precipitate, when produced in the cold, dissolves readily in con-— 
siderable dilute HCl. If it is produced from a boiling solution, basic — 
zirconium sulfate is formed by hydrolysis, which is quite insoluble — 
in dilute HCl (difference from Th and Ce). oe 

10. NasSO4 produces no precipitation, even on boiling the solu- 
tion, which is slightly acid with sulfuric acid (difference from Ti). — 

11. H2O2 precipitates from slightly acid solutions white; volu- 
minous zirconium peroxide, Zr2Os5, which evolves chlorine on being — 
warmed with concentrated HCl. | 

12. NaeS2Oz precipitates zirconium completely as the hydroxide, a 
the precipitate being always contaminated with sulfur. | 

13. Turmeric Paper, after being moistened with the hydrochloric — 
acid solution of a zirconium salt and dried, is colored reddish-brown 
(difference from Th). a 

14. HCl. Ruer * recommends the following test for the identi- — 
cation of zirconium: 





Precipitate the zirconium in the cold by ammonia, filter, wash, and separate 
it from the filter as completely as possible. Dissolve the precipitate in hydro- — 
chloric acid (or if small in amount treat the paper and precipitate together — 
with not too strong HCI and filter.) Evaporate the hydrochloric acid solution — 
to dryness on the water-bath and take up the residue in a little water. To the — 
cold, saturated solution add hydrochloric acid drop by drop, when the presence — 
of zirconium will be evident by the formation of a voluminous precipitate of — 
zirconium oxychloride. Redissolve the precipitate by heating the solution, ~ 
and allow the liquid to cool. After some time fine, silky needles of ZrOCl,.8H,O — 
will precipitate. 

In the somewhat unusual case that zirconium is present in the form of the 4 
insoluble meta-zirconium acid, transform the latter into zirconium sulfate by 3 
heating with concentrated sulfuric acid (2:1), dissolve this in water, precipi- — 
tate the zirconium by ammonia, and carry out the above process. 3 


15. Sodium Iodate produces in slightly acid solutions a volumi- — 
nous white precipitate of zirconium iodate which is soluble in hot, — 
dilute hydrochloric acid (best method of separating zirconium from 
aluminium). } . 

16. Hydrofluoric Acid as a rule spodices no precipitation (differ- ~ 
ence from thorium, cerium and other rare earths). From concentrated 
zirconium solutions a voluminous precipitate may be obtained by the — 
careful addition of hydrofluoric acid, but the precipitate is soluble — 
in an excess of the reagent. | 


* Z. anorg. Chem., 42, 85 (1904). 





METALS OF THE (NH,).8 GROUP 465 


REACTIONS IN THE DRY WAY 


ZrOs is infusible in the oxyhydrogen flame (difference from the 
other earths), but glows brightly. 


THORIUM, Th. At. Wt. 232.4. M. P.>1700<Pt. 


Occurrence.—Thorite (orangite), ThSiO4, with 50 to 58 per 
cent ThO2; thorianite, a mineral discovered in Ceylon, with 72 
to 76 per cent- ThOg and 11 to 12 per cent UO2;* gadolinite, 
Be(Y,Ce,La,Di,Th,O)2FeSiO4, monazite, (Ce,La,Di,Th)PO4, with 2 to 
8 per cent ThO2; and in the rare niobates, samarskite, pyrochlore, 
euxenite, etc. Euxenite is essentially a titanate and niobite of 
Ce(La,Di) and usually contains UOz and FeO. Thorite, monazite, 
and gadolinite are decomposed by acids, preferably sulfuric acid. 


REACTIONS IN THE WET WAY | 


A solution of Th(SO4)2 should be used. 

1. (NH4)OH, (NH4)2S, or KOH produces a white precipitate of 
Th(OH).4, insoluble in an excess of the reagent, but readily soluble 
in dilute acids. By igniting the hydroxide, ThOz is obtained, which 
is soluble in concentrated sulfuric acid only after long digestion. 

2. KeCOz3 or NazCO3 precipitates the white carbonate, soluble in 
an excess of the reagent, and not reprecipitated by the addition of 
ammonia. On boiling, the solution becomes turbid, but clears again 
on cooling. | 

3. (NH2)COxz precipitates the white carbonate, readily soluble in 
an excess; on warming to 50° a basic carbonate is precipitated, 
which redissolves on cooling the solution. Ammonia causes no pre- 
cipitation in this solution. 

4. BaCOs completely precipitates thorium salts in the cold. 

5. KeSOx4 precipitates K4Th(SOs4)4+2H20, difficultly soluble in 
water and insoluble in concentrated K2SO4 solution (difference from 
Y). The corresponding sodium compound is readily soluble in water. 

6. Oxalic Acid precipitates, from solutions which are not too acid, 
all of the thorium as white, crystalline oxalate, practically insoluble 
in oxalic and dilute mineral acids. 

7. Ammonium Oxalate likewise precipitates thorium oxalate, 
which dissolves on boiling with a large excess of this reagent. The 
solution remains clear after cooling, provided the original solution 
did not contain too much free sulfuric acid, and enough ammonium 





* Chem.-Ztg. Rep., 1905, 91. 


466 REACTIONS OF SOME OF THE RARER METALS 


oxalate was used. From the boiling solution of the ammonium 
double oxalate, HCl precipitates practically all of the thorivm as 
oxalate (difference from Zr). 


In the presence of ammonium acetate, ammonium oxalate pro- 


duces no precipitation; but by the addition of HCl almost all of the 
thorium will be precipitated as oxalate. 

8. HF produces a white, gelatinous precipitate, which soon changes 
to a heavy powder. KF causes the same reaction. 

9. Na2S2Oz3 precipitates all of the thorium on boiling. 

There are no characteristic dry reactions. 


SE ee eT ee 


ee ee 


” 
es eo 


_ 7, ‘ 
ve ae ll 


THE GADOLINITE METALS 


YTTRIUM, Y (At. Wt. 89). M.P. between 1000°-1400°, and ERBIUM,* 
Er (At. Wt. 167.7) 


Occurrence.—Yttrium is an important’ constituent of gadolinite, 
Be(Y,Ce,La,Di,Th,Er,O)2FeSiO4, and of yttrotantalite Y(Nb,Ta)Os, 
an isomorphous mixture of yttrium tantalate and yttrium niobate. 
The two elements Y and Er are also found in cerite, thorite, and 
monazite. 

REACTIONS IN THE WET WAY 
A solution of Y(NO3)3 and one of Er(NOs)3 should be used. 
Yttrium Erbium 

1. NH,OH and (NH4)2S pre- Behaves like yttrium. 
cipitate the white hydroxide, in- 
soluble in an excess. 

2. KOH and NaOH precipi- Behaves like yttrium. 
tate the white hydroxide, insoluble 
in an excess; the presence of tar- 
taric acid does not prevent precipi- 
tation; but in this case yttrium 
tartrate is precipitated (difference 
from Al, Be, Th, and Zr). On 
igniting the precipitate, the oxide 
is obtained, which is readily soluble 
in acids. 

3. (NH4)2COs produces a white Behaves like yttrium, except 
precipitate of the carbonate, read- | that the solution does not be- 
ily soluble in an excess of the re- | come turbid on standing. 
agent; after standing some time 
the solution becomes turbid, owing 


to the deposition of a double salt, 
Yo2(COs)3 ° 2(NH4) 2CO3 -2H20. 








*Erbium is not an element itself, but consists of at least three elements— 
holium, thulium, and dysprosium. The separation of these elements is exceedingly 
difficult; for this reason we shall consider simply the reactions of the mixture. 


467 


468 
Yttrium 

4. K2CO3 and NasCOs pre- 
cipitate the white carbonate, read- 
ily soluble in excess; after standing 
some time an insoluble double salt 
separates out. 

5. BaCOgz does not precipitate 
yttrium in the cold, and only in- 
completely on warming. 

6. Oxalic acid precipitates 
white yttrium oxalate, insoluble 
in an excess, difficultly soluble in 
HCl, and perceptibly soluble in 
ammonium oxalate. 

7. K2SO. forms a double salt 
which is soluble in KeSO,4 solu- 
tion (difference from Zr, Th, Ce, 
La, and Di). 

8. HF produces white amor- 
phous YF3, which becomes pulver- 
ulent on warming, and is insoluble 
in water and in HF (difference 
from Al, Be, Ur and Ti). 

Yttrium solutions do not give 
an absorption spectrum. 


REACTIONS IN 


Yttrium oxide is strongly lumi- 
nous on being heated; otherwise 
there is no reaction. 








REACTIONS OF SOME OF THE RARER METALS 


Erbium 


Behaves like yttrium, only the 4 
solution does not become turbid. 
on standing. 


Erbium is not precipitated 
at all, even on warming. 


In erbium solutions, oxalic 
acid produces a light-red, pul- 
verulent precipitate; otherwise 
the reaction is the same as with 
yttrium. 

Behaves like yttrium. 


Erbium solutions give a char- | 
acteristic absorption spectrum:. 


Mu aye bp Me 
683.9 640.5 523.2 468.5 
667.1 548.9 491.6 449.9 
653.5 541.0 487.5 422.3 


649.0 536.4 474.5 416.6 


THE DRY WAY 


Erbium oxide, on being heated 
on a platinum wire, colors the 
flame distinctly green. If the 
light is viewed through a spectro- 
scope, a number of bright lines — 
will be seen in the dark green 
which coincide with the dark 
lines obtained in the absorption — 
spectrum. a 





THE CERITE METALS 


CERIUM, Ce. At. Wt. 140.6, M. Pt. 635°? LANTHANUM, La, At. Wt. 
139.0. M. Pt. 810° 


PRASEODYMIUM, Pr. At. Wt. 140.6. M. Pt. 940°? 


sige an Secale Nd. At. Wt. 144.3. M. Pt. 840°? 


Occurrence.—These three metals are important constituents of 
cerite, H3Ca(Ce,Al)38ig013), and of orthite (allanite), HCazCes3‘ i3013, 
besides being usually found associated with’ the gadolinite earths in 
gadolinite, ete. 


CERIUM 


Cerium forms two oxides, Ce2Os and CeQs ; both are basic an- 
hydrides, from which salts are derived. The cerous salts are white, 
while the ceric salts are orange red. 


REACTIONS IN THE WET WAY 
1. Cerous Salts 


A solution of cerous nitrate, Ce(NO;);, should be used. 


1. NHsOH and (NH4)2S each produce a white precipitate of 
Ce(OH)s3, insoluble in an excess of the reagent, but readily soluble 
in acids. In the presence of tartaric and citric acids, etc., the above 
reagents cause no precipitation (difference from Y). : 

2. NaOH or KOH also precipitate white Ce(OH)s, even in the 
presence of tartaric acid, ete. The white Ce(OH)3 becomes yellow 
gradually on standing in the air, on account of being oxidized to 
Ce(OH)a. 3 

3. KeCO3 and (NH4)2COs each produce a white precipitate insol- 
uble in an excess of the reagent. 

4. Oxalic Acid or Ammonium Oxalate precipitate white cerous 
oxalate, insoluble in an excess of the reagent, and in dilute mineral 
acids. On ignition, pale yellow, insoluble CeOzg is formed.* If the 
oxalate is contaminated with praseodymium oxalate, a cinnamon- 
colored oxide is obtained, which is perfectly soluble in dilute acids. 





* Only when the cerous oxalate is pure. If it contains traces of praseodymium, 
the CeO, is obtained as a bright-yellow powder. 
469 


470 REACTIONS OF SOME OF THE RARER METALS 


5. BaCOs slowly precipitates all of the cerium in the cold. 

6. K,SO4. If a neutral solution of cerous salt is treated with 
solid K2SO4 until no more will dissolve, a crystalline precipitate of 
Ce2(SO4)3-3K2SO4 is slowly formed at room temperature, or more 


quickly on heating. All of the cerium can be precipitated in this way, 


as the double sulfate is insoluble in concentrated potassium sulfate 
solution, but it will dissolve in considerable pure water or more readily 
in acids. From slightly acid solutions cerium can be completely 
precipitated with KeSO4 as Ce2(SO4)3-2K2S04-2H20. (Difference 
from Al, Be and yttrium earths.) 

Na2SOx4 behaves similarly (difference from Th and Zr). 

7. HF produces in neutral and slightly acid solutions of cerous 
salts a gelatinous precipitate which by long digestion with the hot 
solution gradually becomes pulverulent. The precipitate of CeF3-H2O 
is practically insoluble in water and dilute hydrofluoric acid, but readily 
soluble in other mineral acids (difference from Al, Be, Zr and Ti.) 


8. H2O2 colors neutral cerous solutions yellow, but after adding 


an acid the’ color disappears, as the cerium peroxide is reduced to 
cerous salt by hydrogen peroxide and acid. If a cerous salt is treated 
with a slight excess of ammonium hydroxide and then with He2Oo, 

the precipitate becomes reddish orange in color, something like Fe(OH)3. 
_ This is the most sensitive test for cerium; it was discovered by Lecoq 


de Boisbaudran. The composition of the precipitate has been given 


as CeOo-Ce203-HeO2 and as Ce(OH)302H. 


9. Chlorine. If a cerous salt is treated with an excess of alkali — 


hydroxide and then with chlorine gas, a yellow precipitate of 
CeO2-3H20 is obtained. If the chlorine gas is passed through the 
solution for a long time, the precipitate will redissolve. , 

10. Bromine behaves like chlorine except that an excess of this 


oxidizing agent does not dissolve the precipitate. (Difference from 


lanthanum and didymium.) 

11. Cerous salts may be oxidized to ceric salts in acid solutions 
(a) by heating with PbOz and HNOg (1:2); (6) by heating with am- 
monium persulfate; (c) by electrolysis. In all cases, the solution 
_ becomes yellow or orange in color. 


2. Ceric Salts 


A solution of either ceric nitrate, Ce(NO3)4, or of ceric ammonium 
nitrate, Ce(NO3)4-2NH4NO3+H20, should be used. 

The beautiful orange-red color of these solutions is characteristic 
of all ceric salts, as is also their tendency to form difficultly soluble 
basic salts. 





THE CERITE METALS 471 


Preparation of Ceric Compounds.—As has been already stated, cerous 
hydroxide on standing in the air gradually changes to yellow, on account of 
the formation of ceric hydroxide. This oxidation takes place immediately 
on the addition of chlorine or hypochlorites. If the solution of a cerous salt 
is treated with caustic potash solution and chlorine is conducted into it, the 
white cerous hydroxide which was at first formed is quickly changed to light- 
yellow ceric hydroxide. The latter compound dissolves in dilute acids, forming 
orange solutions. It dissolves in concentrated hydrochloric acid with evolution 
of chlorine, forming cerous chloride. If white cerous hydroxide is heated in 
the air, it loses water and is changed into CeO:, which is nearly white when 
cold, dark orange when hot, and is almost entirely insoluble in concentrated 
hydrochloric and nitric acids. In the presence of reducing substances (such 
~as KI, FeSO,, etc.) it dissolves in acids, forming cerous salts: 


2CeO.+8HCl+ 2KI =2KCl+4H,0+1.+2CeCh. 


CeO, can also be dissolved by warming with concentrated sulfuric acid, 
with evolution of oxygen and formation of cerous sulfate. It can be readily 
brought into solution by fusing with potassium pyrosulfate and dissolving 
the melt in considerable hot water to which a little acid is added. 

If a mixture of cerous and praseodymium hydroxides is ignited in the air, 
a cinnamon-colored mass is obtained, which contains all of the cerium as dioxide 
and is readily soluble in dilute acids, forming ceric salts. If concentrated 
HCl is used, there is an evolution of chlorine, the ceric salt being reduced to 
cerous chloride. Concentrated nitric acid dissolves it, forming cerous and ceric 
salts; a distinct evolution of oxygen can always be detected. 

The reason why the brown mass containing a little praseodymium dissolves 
although the pure oxide does not, is probably the following: CeOkz, like MnO, 
and PbO, (see pp. 162 and 206), plays the part of an acid anhydride, so that 
the brown mass contains the praseodymium as the salt of ceric acid. On 
treating this salt with a stronger acid, the praseodymium salt of the latter is 
formed, setting free ceric acid (ceric hydroxide), which in the hydrated form 
is readily soluble in acids, forming ceric salts. 

Basic Ceric Salts.—Ifa solution of ceric nitrate is evaporated on the water- 
bath to a consistency of syrup, the mass dissolves readily in water after it has 
become cold, and the solution can be boiled without becoming turbid. If, 
however, a little nitric acid is added, a yellow precipitate is immediately formed, 
consisting of basic ceric nitrate; on the addition of more acid the precipitate 
redissolves. This can be explained as follows: By treating the solution of 
ceric nitrate with considerable water it becomes hydrolyzed considerably, but 
the basic salt produced is present in the hydrosole state and is changed by the 
acid into the hydrogele form. 

As lanthanum and didymium salts do not yield basic salts under these con- 
ditions, this property can be used for separating cerium from these metals. 

It is characteristic of cerium to form with ammonium nitrate an easily 
crystallizable salt, ceric ammonium nitrate: Ce(NQO;),-2NH,NO;-H,0O. 

All ceric salts may be readily reduced by the ordinary reducing ses 
(aleohol, HI, SO2, H:S, HNO2, H2O:, ete. ) to cerous salts. 


Oxalic Acid added to a concentrated solution of a ceric salt at 
first precipitates a dirty orange precipitate which gradually becomes 


; 


472 REACTIONS OF SOME OF THE RARER METALS 


yellow and gelatinous as the addition of oxalic acid is continued, and — 
finally crystalline. The precipitate dissolves in a large excess of oxalic 
acid, but the solution gradually becomes turbid in the cold, or more — 
quickly on heating, as the ceric salt is reduced to cerous salt at the — 
expense of the oxalic acid from which carbon dioxide is evolved. Finally 
all the cerium will be precipitated as cerous oxalate. (Difference from 
La, Di and the yttrium earths.) ; 
Ammonium Oxalate behaves similarly. 


REACTIONS IN THE DRY WAY 


The borax head is colored dark brown when hot and light yellow 
to colorless when cold, after being heated in the oxidizing flame. In — 
the reducing flame the bead becomes colorless, although strongly 
ignited CeOz will remain suspended in the bead, giving it a Bi 
yellowish appearance. 


LANTHANUM, La. At. Wt. 138.9. M. Pt. 810? 


Lanthanum forms only one oxide,* Laz03, which, even after being _ 
strongly ignited, dissolves readily in acids. Its salts are colorless and 


yield no absorption spectrum, so that lanthanum may be distinguished 


in this way from didymium and erbium. 


REACTIONS IN THE WET WAY 


A solution of lanthanum nitrate, La(NO3)3, should be used. + 

1. NH:OH and (NH4)2S precipitate a white basic salt which is 

difficult to filter. The presence of tartaric acid prevents the precipi- 
tation. (Difference from yttrium.) 

2. KOH and NaOH precipitate the white hydroxide, La(OH)3. 
There is no change to be noticed on treating with oxidizing agents 
(difference from Ce). La(OH)3 is soluble enough in water to turn 
red litmus-paper blue, and it decomposes ammonium salts on warm- 
ing with evolution of ammonia. The fused oxide is realy soluble | 
in car 

3. (NH4)2CO3 produces a , white vrei slightly soluble in an — 
excess of the reagent (difference from aluminium); after standing some 
time crystalline lanthanum ammonium carbonate is precipitated. 

4. Oxalic Acid produces a white crystalline precipitate, insoluble 
an excess of the precipitant and in ammonium oxalate, but soluble in 
dilute mineral acids. (Difference from Th.) . 





*H.O. is said to cause the formation of LazQ;. Cf. Z. anorg. Chem., 21, 
70 (1899). 





THE CERITE METALS 473 


5. KoSO. precipitates white, crystalline Laz(SO4)3-3K28O4, in- 
soluble in a concentrated KeSO, solution. 

6. Lanthanum Sulfate is soluble only in ice-cold water; on warm- | 
ing the saturated solution to 30° the salt separates out thickly (differ- 
ence from cerium.) 

7. Iodine —If ammonia is added to a cold, dilute acetic acid 
solution of a lanthanum salt, and the slimy precipitate is washed with 
water and then treated with solid iodine, the whole mass gradually 
assumes a blue color which is similar to that produced by the action 
of iodine upon starch (this property is peculiar to lanthanum). The 
blue color is destroyed by the addition of acids or alkalies, and dis- 
tinguishes it from all the other earths. 

8. HF precipitates white gelatinous lanthanum fluoride, which 
eventually becomes crystalline LaF3-3H2O; the precipitate is insoluble 
in an excess of the precipitant and in dilute acid it is gradually dissolved 
by strong mineral acids. 


NEODYMIUM, Nd. At. Wt. 143.6. M. Pt. 840°? 


yt icgedae PRASEODYMIUM, Pr. At. Wt. 140.5. M. Pt. 940°? 


It is very difficult to separate these two metals from one another. 
It is accomplished only by repeated fractional crystallization of the 
ammonium double nitrates. 

Neodymium apparently forms only one oxide, _Nd2Qz; _ it 
appears bluish after being ignited, and is readily soluble in acids, 
forming violet salts, which afford a characteristic absorption spec- 
trum. 

Praseodymium, on the other hand, forms a greenish-white oxide, 
Pr2O3, which on being ignited is changed into dark-brown peroxide, 
Pr4O7. On being heated in a stream of hydrogen, the latter is reduced 
back to Pr2O3. The peroxide dissolves in acids with loss of oxygen, 
forming greenish salts corresponding to the lower oxide and yielding 
a characteristic absorption spectrum. 

The Didymium reactions take place with a mixture of the two 
elements. A solution of didymium nitrate, Di(NOs)s, is used. 

Didymium salts are violet and show a characteristic absorp-. 
tion spectrum (difference from Ce and La). The behavior toward 
NH4OH, (NH4)2S, KOH, (NH4)2COsz, and KeSOs is exactly the same 
as with lanthanum. Oxalic acid precipitates the reddish sige 
which in other respects is like lanthanum oxalate. 

Absorption Spectrum of Praseodymium.—596.9up, 5904yu, 482.1upy, 
469.5uyu, and 444.1yu. (Continued on p. 480.) 


REACTIONS OF SOME OF THE RARER METALS 


474 


‘pour1oy St yes e[qnop 


*sso0Xe UI o[qnyos 


*peulioy wayyy SI 0981} 3 
-184 UIntI434 qnq ‘aon 
-eyidioeid eq} quaaoid 

















- ‘ a[qnjOsUI UB sInoYy euos 
SESE ON "OO" WtM Lee IyFY  "sseoxe Ur o9iq | -UT oyezIdwoord 714 yOU SeOp plo OLIe4zIe TL, *O*A 
-jos 97841d1001d on *sS90xe UI 9[qnyjos 
-UI 93ezIdro0I1d eazy M 
*Burjooo 
Sc inr aati: ties “s[BJOUI o}1100 9Y} 910J | “UOT;eZIdIOeId ayy syUAA 
*BulIog UO peyeyzidio . . -0q peyeqrdioeid st -1d poe 9a 
sed sjusdoo wouy | SOON HM sw sosUTPRE | LL, HONAN, Jo, won|“ PMooxa wy aignyos | Nasooxo tr apgnjos | "OWL 
you pue ‘ssvoxa ur eqn “UI oyeyIdiosid ay “Ur ayeyidioeid ayy 
-[os o7Byidoord O7TY MA 
San “Bul "HO'HN 4qG ‘uoT}BYIdIeId 94} sJUVA 
uo (anjjns yzTM ert -[Ioq uo payeqdioider | poyezidioeider ynq ‘sseo *sso0xo0 UI efqnjos | -o1d §=ploe  o1eqiey t 
24% ra 19 oe A AA qnq ‘ssooxe e818] UI sjqn | -xo Ul aTqnjos ApZYysYs | -ur oyeqrdiooid o7I MA *sseoxe UI eyqnyos OFZ 
FERRE mt -[Os o3eyidiooid a4IG MM 27 B4Id100I ov MA. “Ul oyByIdiosid ofIy 
‘uoreyidioeid oy} Ss}UOA 
uo pezeqid he tree bajagscgole ties: sse0xe UI e[qnyjos eM cS aaa. 
‘azey1dioeid 0 Tor B Ul 9IqNTos AP~NoWT ; at 
yet N Sse0Xe UT a[qnjos ATIpwal | 67, alea ts 1d ae ing *P “ur ayeqidieid ayy A -yn[Ip 10 Burro wo pazey ord 
eyeyridroeid ov1GUM SP , -1d1dei1del ‘sseoxe UI eyqn 
~[Os a7eyidiooid ey1Iy 
“WO|NIOS [214 ‘uot ByIdioeid 8yy 
"nou ysourye ot} (og *Ss00xX9 UI 9]qnyjos *sSa90xe UI a7qnyjos *SS90xo UI o[qnyos | syUeAoId plow o1lezIe T, <OTy 


uo ozeyidiweid o71IG MK 


-ul eyeqidioeid OTT M 


-ur opezyidiooid o7IY 


“ut oyeqidiooid ayy A 


*ssooxe UL eqn 
-[0s oyeqidroeid o71 A 





*O*S°eN 





*00*(HN) 





**OO*%N 10 
OO 





S°(°HIN) 10 
HO'HN 





HO®N 10 
HOM 








SHLYUVA AHL JO SNOILOVaEY AHL JO AYVWWOAS 


475 


THE CERITE METALS 


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log UG *sSe0xe UI 8[qn 
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REACTIONS OF SOME OF THE RARER METALS 





476 


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snoydaiows ‘ayy A : 
1OH 14 . 
-BlopIsu0d Aq peje} 
-dmeid si worjnjos 
: oy.L ‘plow sey Yonw ‘ayBjooe UINTUOUL 
MOT NLOS "OS \ Acob amine duanaautos 00} UIe}U0D you Pprp | -UIe UL eTqnyos ‘spre 
peyerueoug) = @ UW “aUON PI a is 4 ie bilge ‘ayeqidiooid ON UONOS [VUISIIO oY} | [RIOUTUT OINTIp UI pus | %;YAL 
etqnjosur sha eqqnop WS st Macs Sy ba 4eq} peptaoid Sut| sseoxe ur afqnjosut 
Bi: SAR ene -]009 uo pazeyzidroord | ozeztdtoord «ORT AA 
-91 you pue ‘suT[LOog 
uo sseoxe UT e[qnjos 
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‘ITOH 
ul 9Tqnjosur 4sour 
-[@ SI yt UoyM ‘UOT} . 
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peyeyidiord = SBAL Ris oop oe -Jopisuoo jo uorpppe | oynpIp AreA ut afqnzos 
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OSH aeepasiontee ‘OOrE JH ¥0%*("HN) PY exOQ 























(‘panuyuo))—SHLUVA GAUVU AHL AO SNOILOVaY HHL JO AUVWNOAS 


477 


THE CERITE METALS 


‘O68T ‘Bunjnaz~wayD ‘IAISVID *) Aq poystiqnd osoy} jo djay oy} 441M popidulod o19M solqe} osey [—' ALON 





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SI 388 e[qnop VW 


‘OIystIopOVIVYO 


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snoulzejes = 4olOlA 


' *ploe o1exo 


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SMOYS 98S PI]OS OYL 
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st 978s ejqnop y 


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snounysjes § oayIUM 


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‘ayeqyidioeid snou 
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St 988 e[qnop V 


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“pyoo oy} Ur 
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-woo ynq ATMOTE 


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snourzees = sayIG MA 


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oyeyIdioeid az MM 


“Sploe [B1OUTUL 
eynTIp Ul puB sprloB 
oyjVxo Ul 
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erqnyosur 


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‘OIYSTIOPOVIVYO 


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pezyeyidioeid = Ajayoyd 
-mmoo0ur ‘plod oy} Ur 
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‘ayveyidiosid snou 
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REACTIONS OF SOME OF THE RARER METALS 


‘dnois g*H 94} 
JO s[ejour oy} Jo ows uoqjo pus ‘O*H ‘O*EN ‘O3W ‘OPA ‘O° “OA “OWT “OIC *O%D “ONT “O%A “OIS :survzUOO opUTPOpEH 


(ALIYHO) ALINITOCVS AO SISATVNY 


478 


479 


THE CERITE METALS 


“£8 TENSN OY} UI IO} Pose} OIE UINTUTUIN[S pues UOIT y 








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fo aouasaid ay}2 smoys a7nj0 
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fo uoyouwof ay? “JH 





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*(HO)20 :ALVLIdIOGUg 





480 REACTIONS OF SOME OF THE RARER METALS 
Absorption Spectrum of Neodymium.—T29.1yp, 690.6up, (579.7up, 


675.9uu), 531.7up, (522.2uu, 520.9uu), 512.0uu, 509.6uy, 482.1yy, 
475.9up, 469.5uu, 461 4up, 444.3, 434.1lup, 427.7uy, 417.3up. 


TANTALUM, Ta. At. Wt. 181, and NIOBIUM, Nb. At. Wt. 93.5 


These two rare elements, belonging to the nitrogen-vanadium group, 
form oxides of the formula R2Os5, which behave as acid anhydrides 


and should have been considered, perhaps, under the acids. Since, — 


however, tantalic and niobic acids are soluble, under certain conditions, 


in strong acids and from these solutions are precipitated by ammonia — 


and ammonium sulphide, it seems better to consider them at this place. 

Occurrence.—In the form of meta-acids these elements appear in 
the isomorphous minerals tantalite, Fe(TaOg3)2 and niobite or columbite, 
Fe(NbO3)2. In tantalite a part of the tantalic acid is replaced by 
niobic acid and a part of the iron by manganese. Niobite shows an 
analogous behavior. 

In the form of pyro acids the two elements occur as an isomorphous 
mixture in the mineral yttrotantalite, Y4(Ta2O7)3 and Y4(Nb207)3. 
Finally tantalum, and to some extent niobium, replaces the phos- 
phorous’ in monazite; (Ce,La,Di)POs. Tin is usually found in all 


the above minerals and often tungsten; conversely cassiterite and — 


wolframite often contain small quantities of niobic and tantalic acids. 


TANTALUM, Ta. At. Wt. 181. M. Pt.=2850°. Sp. Gr.=16.5 


Metallic tantalum* is ductile, although the presence of a little 
impurity makes it harder than tool steel. On ignition in the air it — 


assumes a yellow to blue tinge caused by a thin coating of oxide. 
Tantalum is not attacked by boiling H2SO4, HCl, HNOs, or even aqua 


regia, but it is slowly dissolved by hydrofluoric acid with evolution — 
of hydrogen; any metal remaining undissolved is then brittle on — 
account of absorbed hydrogen. The concentrated solution in hydro- — 
fluoric acid forms with concentrated KOH insoluble, crystalline potas- — 
sium tantalum fluoride, KoTaF7. By evaporating the solution in — 
HF with concentrated H2SO,4 until the former acid is all expelled, the — 
residue dissolves in a little cold water, but the solution becnaas turbid | 


on dilution or especially by boiling. 
Tantalum forms two oxides, Ta2O4 and Ta2QOs, the former being 


indifferent chemically and the latter an acid anhydride. After ignition, — 
the pentoxide is insoluble in acid and is not rendered soluble by fusion — 





*W. von Botton, Z. Electrochemie, 11, 45 (1905). 


7 << wre 





THE CERITE METALS 481 


with pyrosulfate, although it is volatilized by heating with ammonium 
fluoride. Fusion of the oxide with caustic alkali in a silver crucible 
gives rise to alkali tantalates, both the meta- and hexatantalates being 
known; only the former are soluble in water. 

Potassium hexatantalate, KeTagO19+16H20, is soluble in water 
and caustic potash solution, while the sodium salt is soluble in water, 
but not in caustic soda. If potassium hexatantalate is treated with 
hot water, a part of it goes into solution and undergoes hydrolysis, 
forming a colloidal solution of HgsTagO19. If COe is conducted into the 
solution, the tantalic acid is completely precipitated in a flocculent 
condition. (Difference from niobic acid.) The remaining tantalates 
are all insoluble. 


REACTIONS IN THE WET WAY 


A solution of potassium hexatantalate should be used. 

1. Mineral Acids.—(a) H2SO.4 precipitates tantalic acid from 
cold, dilute solutions, and the precipitation becomes nearly quantita- 
tive on boiling. Hot, concentrated H2SOx4 dissolves the precipitate 
produced by the dilute acid. On diluting the solution with water 
after it has become cold, the tantalic acid is reprecipitated (difference 
from niobium.) 

(b) HCl added to a concentrated solution at first produces a pre- 
cipitate, which dissolves in an excess of the acid, forming an opalescent 
solution. From this solution sulfuric acid precipitates tantalic acid 
in the cold, but the precipitation is not quantitative even on boiling. 

(c) HNOs has the same action as HCl. 

2. NH4OH and (NHsgz)2S precipitate from the hydrochloric acid 
solution either tantalic acid itself or an acid ammonium tantalate; 
tartaric acid prevents the precipitation. 

3. NaeCOz produces a .partial precipitation of tantalic acid when 
added to an acid solution of a tantalate, but the precipitate dissolves 
in an excess of the precipitant. 

4. HeS precipitates tantalic acid in the cold, especially from sul- 
furic acid solution, and the precipitate is almost quantitative. 

5. H2O2 dissolves freshly precipitated tantalic acid if acid or 
alkali is present. From the solution thus obtained, the tantalic acid 
is not precipitated by the above reagents unless the hydrogen peroxide 
is destroyed by boiling the alkaline solution or by the action of sulfurous 
acid. 

6. Ks[Fe(CN)6| and KCNS produce white precipitates. 

7. Tincture of Nutgalls produces no precipitate (difference from 
niobic acid). | 


482 REACTIONS OF SOME OF THE RARER METALS 


8. K4[Fe(CN).] produces in acid solutions a light-yellow precipi- — 
.. tate, which becomes brown on the addition of a little ammonia. 


9. H{KF>2|.—If a concentrated solution of tantalic acid in hydro- 
fluoric acid is treated with KF, the difficultly soluble Keo{TaFz]. is 
formed, which separates from the solution in the form of orthorhombic 


needles (200 parts of water dissolve 1 part of salt) difference from 


niobium). On boiling the solution of tantalic potassium fluoride, the 
very difficultly soluble oxyfluoride precipitates (K4TasOsFi4). By 
means of this reaction the merest trace of tantalic acid can be detected 
in the presence of niobic acid. 

10. Zn and HCl do not produce colored solutions (difference from 
niobium). 

REACTIONS IN THE DRY WAY 

Ta2Os5 is infusible. The bead of salt of phosphorus remains 

colorless in both oxidizing and reducing flames. The addition of 


FeSO. does not cause the formation of a blood-red color. (Difference 
from Ti and Nb.) 


NIoBIUM. Sp. Gr.=12.7. M. Pt.=1950° C.* 


The metal niobium is very similar to tantalum; it is more readily 


attacked by acids. 

Niobium forms three oxides: Nb2O2, Nb204, and Nb2Os, of which 
the last is an acid anhydride. Nb2Qs, like Ta2Qs, is insoluble in acids 
after it has been ignited and is not rendered soluble by fusing with 
potassium pyrosulfate. The melt dissolves in cold water, but niobic 
acid separates from the solution on boiling. By fusing with KOH or 
KeCO3, potassium hexaniobate, KgNbeg0i9:16H20, is formed, which 
is soluble in water. The corresponding sodium salt is insoluble in 
caustic soda solution, but soluble in water. 

Sodium hexaniobate is largely hydrolyzed in aqueous solution, form- 
ing a colloidal solution of niobic acid, HgNbeOi9. By passing CO2 
through this solution for half an hour, there is no precipitation of 
niobic acid, but only after long standing. (Difference from tantalic 
acid.) 

REACTIONS IN THE WET WAY 


Use a solution of potassium hexaniobate. 

1. Mineral Acids produce in alkali niobate solutions a white, 
amorphous precipitate of niobic acid, which is only slightly soluble 
in an excess of the acid. Concentrated sulfuric acid, however, dissolves 





* WERNER VON Bouton, Chem. Zentralbl., 1905, I, 586. 








re ae 


THE CERITE METALS 483 


the niobie acid on warming; and the solution remains clear after being 
diluted with cold water. (Difference from tantalum.) 

By boiling the diluted solution, the niobic acid is almost com- 
pletely precipitated, but in a very finely divided condition, such that 
it is hard to filter. If the acid is exactly neutralized with ammonia, 
the niobic acid is precipitated in a flocculent condition easy to filter. 
It is best to wash such a precipitate with 0.5 per cent ammonia water 
or with 1 per cent acetic acid. Washing with pure water causes a 
turbid filtrate and mineral acids should not be used. 

If a solution of niobic acid in sulfuric acid is poured into a con- 
centrated solution of ammonium sulfate, no niobic acid is precipitated 
by boiling. (Difference from tantalic acid.) 

If the niobic acid is treated with boiling hydrochloric acid, it dis- 
solves only slightly, but on pouring off the acid, the residue is soluble 
in water.* 

Carbonic acid decomposes sodium niobate to some extent. On 
the other hand, a niobate solution prepared after fusing with sodium 
carbonate and potassium nitrate is not decomposed by COe. (Differ- 
ence from tantalic acid.) 

2. NHiOH and (NH4)2S precipitate niobic acid from the sul- 
phuric acid solution, and the precipitate is soluble in HF. 

3. H2O2 reacts as with a tantalate. 

4. Tincture of Nutgalls produces no precipitate. 

5. K4{Fe(CN).] produces a grayish-green precipitate. 

6. H{[KF2].—If niobic acid is dissolved in an excess of HF and 
KF is then added, readily soluble niobic potassium fluoride is formed 
(12.5 parts of water dissolve 1 part of the salt). By boiling the dilute 
aqueous solution, soluble potassium niobic oxyfluoride is formed, which 
is even more soluble. (Difference from tantalum.) 

7. Zine produces in an acid solution of a niobate a dirty-blue color- 
ation which disappears after some time. (Difference from tantalum.) 


REACTIONS IN THE DRY WAY 


The bead of salt of phosphorus is blue, violet, or brown in the 
reducing flame (according to the amount of niobic acid which is 
present); the bead becomes red on the addition of FeSOs. 





* This behavior reminds one of metastannic acid. (Cf. p. 257.) 


484 REACTIONS OF SOME OF THE RARER METALS 


Separation of Tantalum from Niobium a 


* ey 
ae, 


According to Weiss-Landecker* and Hauser-Lewite.t | =a 





Fuse the two oxides with a little sodium carbonate in a platinum crucible 
cool, extract with hot water and filter off the undissolved sodium hexatanalate. — 
Wash the residue with a solution of NaHCO; and conduct CO; into the filtrate, — 
whereby some flocculent tantalic acid will be precipitated, but the niobic acid — 
will remain in solution. Filter, unite this precipitate with the residue from — 
the sodium carbonate fusion and dissolve both in sulfuric acid and hydrogen — 4 
peroxide. Pass SO, into this solution and boil; tantalic acid will be precipi- — 
tated. Similarly saturate the sodium carbonate solution of sodium niobate — 
with SO, and niobic acid will be precipitated by boiling. aa 





- ey 
Vig 


*Z. anorg. Chem., 64, 65-103 (1909). 
7 Z. angew. Chem., 1912, 100. 


METALS OF THE H2S GROUP 


THALLIUM, VANADIUM, MOLYBDENUM, TUNGSTEN, SELENIUM, 
TELLURIUM, RHODIUM, PALLADIUM, OSMIUM, IRIDIUM, RU-. 
THENIUM 


THALLIUM, Tl. At. Wt. 204.0. Sp. Gr.=11.9. M. Pt.=302° 


Occurrence.—Thallium is found in nature very sparingly; in small 
amount in many varieties of pyrite, and accompanying potassium in 
carnalite and sylvite, in many lithium micas and in many mineral waters. 
It replaces the silver to a considerable extent in copper-silver selenide, 
in crookesite, (AgT]Cu)2Se, and in berzelianite, (CuAgTl)2Se. There 
are no characteristic thallium minerals. The principal sources of our 
thallium is the dust from sulfuric acid plants where pyrite containing 
thallium is used. 

Metallic thallium reminds one of lead in its color, softness, high 
specific gravity, and low melting-point. 

Thallium dissolves readily in nitric and sulfuric acids, but not in 
hydrochloric acid. It forms two oxides: thallous oxide, Tl2O, and 
thallic oxide, Tl,03; both are anhydrides of bases and from them 
thallous and thallic salts are derived. 


REACTIONS IN THE WET WAY 


A. Thallous Compounds 


Thallous compounds are colorless and soluble in water as a rule. The sul- 
fide, chloride, bromide, iodide, and chromate are insoluble in water. Thallous 
oxide is a colorless powder, whose aqueous solution reacts alkaline and absorbs 
carbon dioxide with avidity. 

Use a solution of thallous sulfate for the following reactions: 


1. HeS causes no precipitation from solutions which contain 
mineral acids; in neutral solutions, thallium is incompletely precipi- 
tated as black thallous sulfide, TleS. ToS is readily soluble in mineral 
acids, but insoluble in acetic acid and alkaline sulfides. It is oxidized 
readily on standing in the air to thallous sulfate. 

2. (NH4)2S precipitates all of the thallium as TlS. 

3. KOH, NaOH, or NH4OH produces no precipitation. 

485 


486 REACTIONS OF SOME OF THE RARER METALS 


4, Alkali Carbonates cause precipitation only in very concen- 
trated solutions, for thallous carbonate is fairly soluble (100 parts of 
water dissolve 5 parts of the salt.) 


5. HCl produces a heavy, white precipitate of thallous chloride, — 


very slightly soluble in water, and still less so in water containing a 
little hydrochloric acid. 7 
6. KI precipitates yellow thallous iodide, Tl, from even the most — 
dilute solutions; this is the most sensitive reaction for thallium. 
7. Alkali Chromates precipitate yellow thallous chromate, insoluble 
in cold nitric or sulfuric acids. 

8. Ho[PtCle] precipitates light-yellow thallium chloroplatinate, 
which is almost entirely insoluble in water; 1 part dissolves in 15,600 
parts of water at 15° C. and in 1950 parts of water at 100° C. 

9. Alo(SO4)3.—If a solution of thallous sulfate is treated with 


aluminium sulfate and the solution is then allowed to crystal- 4 
lize, glistening, colorless octahedrons are obtained of thallium alum, 


TIAI(SO4)2+12H20. | 
10. K3[Fe(CN)6] precipitates brown T1(OH)s3 in alkaline solutions: 


2[Fe(CN)¢|=+30H~+TI* — 2[Fe(CN)6]=+T1(OH)s. 


Thallium is like lead in respect tc its specific gravity and to the solubility 


of its halogen compounds; but, on the other hand, it is similar to the alkalies 
with regard to the solubility and alkaline reaction of the hydroxide and car-— 
bonate, and with regard to its forming an insoluble chloroplatinate and an 
alum. 


B. Thallic Compounds 


Thallic compounds cannot as a rule be prepared by the oxidation of thallous 
compounds (with the exception of thallic chloride, which is readily obtained 
by the action of chlorine water upon thallous chloride). They are obtained 
by the solution of thallic oxide * in acids, and can be distinguished from thallous 
compounds by the ease with which they suffer decomposition in aqueous 
solution. Thus thallic sulfate is decomposed on boiling its aqueous solution 
into thallic hydroxide and sulfuric acid; the nitrate behaves similarly. 

The chloride, TICl;, is a hygroscopic and not very stable substance; on 
being heated to 100° C. chlorine is evolved with the formation of thallous 
chloride. 


1. KOH, NaOH, and NH,OH8 precipitate brown thallic hydroxide, 
TI(OH)3, from solutions of thallic salts, which changes to TIO(OH) 
on standing in the air, is difficultly soluble in acids and insoluble in 
an excess of alkali. 


- 





* T1.O; is not attacked in the cold by concentrated sulfuric acid, but is dissolved 
on warming. The hydrated oxide, TIO(OH), is much more soluble. 





METALS OF THE H.S GROUP 487 


2. HCl and alkali chromates do not cause precipitation. 
3. KI precipitates thallous iodide with deposition of iodine. 


REACTIONS IN THE DRY WAY 


Thallium salts color the non-luminous gas-flame a beautiful emerald 
green. The thallium spectrum consists of a green line at 535.0, nearly 
coincident with the green barium line at 534.7. 


VANADIUM, V. At. Wt. 51.0. M. Pt.=1730° 


Occurrence.—Vanadinite, Pbs(VO4)3Cl; carnotite; * mottramite, 
(CuPb)2V2010+2H20; many clays and in almost all granites. 

| Vanadium, like nitrogen, forms five oxides: V20, V2O2, V20s, V2O0a, 

V20s. 

The first three of these oxides are basic anhydrides. Compounds 
representing these valencies of vanadium are not encountered in quali- 
tative analysis except, to some extent, in the tests for vanadium with 
strong reducing agents. The oxides V2O4 and V20Os represent the types 
of vanadium compounds usually encountered in analytical chemistry. 

V204 is the anhydride of hypovanadic acid, V202(OH)4. This com- 
pound is an amphoteric substance and forms salts with both acids and 
bases. V2O« itself is a blue powder, soluble in concentrated acids, 
forming blue divanady] salts: 


V204+2H2S04 = V202(S04)2+2H20. 


If the solution of divanadyl] sulfate is treated with sodium carbonate 
or ammonia (avoiding excess), hypovanadic acid separates out as a 
grayish-white precipitate, which, like the anhydride, is soluble in acids 
with blue color and in alkalies with a brown color. The alkali hypo- 
vanadates correspond to the symbols NazV205 and NazV4O9. Hypo- 
vanadates of other metals are for the most part insoluble in water and 
such precipitates may form when an acid solution containing vanadium 
is neutralized. The divanadyl compounds are readily formed by 
reducing solutions of the pentoxide in mineral acids with sulfurous acid 
(cf. p. 489), and serve, on account of their blue color, for the detec- 
tion of vanadium. 

V20s is the anhydride of vanadic acid and is an orange-red crystal- 
line mass, which is readily fusible but non-volatile. It is only slightly 





* According to FrrepeL and CuMENGE, carnotite contains 18 per cent V2O; and 
55 per cent VOs, as well as K, Ca, Ba, H, As, and P. (Cf. HiLLEBRAND and Ran- 
soME, Am. J. Science, 10, 138.) 


488 REACTIONS OF SOME OF THE RARER METALS 


soluble in water; forming a slightly acid, yellow solution, but readily — 
soluble in concentrated solutions of caustic alkalies, forming vanadates. — 
Like phosphoric acid, vanadic acid exists in the form of meta-, 
pyro-, ortho-, and poly-compounds, of which the meta-compounds are ~ 
the most stable and the ortho-compounds the least so. Thus an aqueous 7 
solution of potassium or sodium orthovanadate is hydrolyzed, even in — 
the cold, into the pyro-salt and alkali hydroxide, 


2Na3VOs+H20 @ NasV207+2Na0H, 
and on boiling the meta-salt is formed: 
Na4V207+ H20 @ 2NaVO3+2Na0H. 


The meta-, pyro-, and ortho-salts of the alkalies are colorless or 
slightly yellow, while the polyvanadates, e.g., the tetra- and hexa-— a 
vanadates, are intensely orange or reddish. Thus the colorless or — 
light-yellow solutions of the ortho-, meta-, and pyrovanadates are — 
colored intensely orange on the addition of acid. 

Besides the above types of vanadium compounds, this element 4 
exists as pervanadic acid, HVO,, formed by the addition of hydrogen — 4 
peroxide to the acid solution of a vanadate. 3 

The reactions of quadrivalent and quinquevalent vanadium will be 
considered together; the other forms are not common enough to make” 7 
it necessary to describe their characteristic reactions. . 


REACTIONS IN THE WET WAY r 


1. NH,Cl.—If a piece of solid ammonium chloride is added to a — 
solution of an alkali vanadate, colorless ammonium metavanadate — 
separates out, 


NasV207+4N H4Cl=2NH4V03+2NH3+H20+4NaCl, 


difficultly soluble in a concentrated solution of ammonium chloride. 

2. Pb(C2H302)2 precipitates vanadic acid quantitatively as vheme zs 
lead vanadate: . 
3Pbt *+2VO04 > Pb3(VO4)e. 


This precipitate, however, is more soluble in dilute nitric acid than is lead i 
chromate, and it is possible to separate chromic acid from vanadiec acid by — 
treating the solution of the two acids with lead nitrate in dilute nitric acid 4 
solution; under the proper conditions all but a fraction of a milligram of the 
chromium is precipitated and 100 mgms. of vanadium yield no precipitate.* — 
The vanadium can be detected in the filtrate by the HO: test. a 


* The solution is neutralized exactly with NaOH, 2 cc. of HNO; (sp.gr. 1 20) 
are added together with an excess of 20 per cent Pb(NOs)2 solution. 








METALS OF THE HS GROUP 489 


3. NH,OH added to a solution of an alkali vanadate causes no 
precipitation. Vanadic acid behaves like phosphoric acid toward 
ammonia; vanadates of ferric iron, aluminium and uranium are likely 
to be precipitated by ammonia, as well as vanadates of the alkaline 
earth metals. 

NH4OH added to a solution of a vanadyl salt precipitates dark- 
gray hypovanadic acid: 


V202Cl4,+4NH40H =4N H4Cl4+ V202(OH)a. 


The precipitation is not quantitative and small quantities of vanadium 
may remain in solution when the vanadyl salt is alone present. If, however, 
an excess of ferric chloride is added to the solution, the vanadium is quan- 
titatively precipitated upon the addition of ammonia. This is true both of 
vanadie acid and of vanadyl salt, ferric vanadate and ferric hypovanadate 
being precipitated with the ferric hydroxide. Instead of ferric chloride, an 
aluminium or uranium salt may be used for the same purpose. 


4. (NH4)2S produces no precipitation, but causes the solution 
to turn brown, owing to the formation of thio-salts. 


If hydrogen sulfide is conducted into a strongly ammoniacal solution of 
a vanadate or hypovanadate, the solution at first turns yellowish-red, but the - 
color slowly deepens until eventually a characteristic, brilliant, violet-red color 
is obtained when the solution has become saturated with H.S. Ammonium 
salts interfere somewhat with this test, but their influence is overcome by the 
addition of a large excess of ammonia. The red color is probably caused by 
the formation of ammonium thiovanadate. As little as 0.2 mgm. of vanadium 
can be recognized by the red color. 

The addition of acid to the red solution produces a black precipitate of 
V8, or V.S;. The precipitation is not quantitative; the filtrate is always 
colored blue and contains detectable amounts of vanadyl salts. The precipi- 
tate is soluble in alkalies, alkali carbonates, and in alkali sulfides, forming a 
brown solution. 

Molybdenum gives a similar red color, in case it was not completely removed 
by previous treatment with hydrogen sulfide in acid solution, and obscures the 
above test or may be mistaken for vanadium. 


5. H2S gives no precipitation in acid solution, but reduces com- 
pounds of vanadic acid to divanadyl compounds, so that the solution is 
colored blue: 


2H3VO4+H28+4HCl= V202Cl4+6H20-45. 


6. Reducing Agents (SO2, H2S, HBr, alcohol, oxalic and tartaric 
acids, sugar, etc.) reduce acid solutions containing vanadates to blue 
vanadyl salts: 


pee ee 
2V04=+S037 +10H*t —> [V202]* +4S$0,-+5H.20. 


490 REACTIONS OF SOME OF THE RARER METALS 


HI reduces vanadic acid to green salt of V20s3: 
V0O,*+21-+-8H* — Vit t4To +4H.20. 


The green color only appears after the iodine has been removed by — 
continued boiling of the solution. 


Metals, such as Zn, Al, and Cd, cause still further reduction of vanadie 
acid, so that the solution turns at first blue, then green, and finally violet. j 

Boiling an acid solution of a vanadate with concentrated hydrochloric acid 
and alcohol reduces the vanadium quantitatively to divanadyl salt. Treat- 
ment of vanadic acid with ferrous salt also reduces the former to divanadyl — 
salt and the excess of the ferrous iron can be oxidized by cold potassium di- — 4 
chromate solution without oxidizing the vanadium. a 


7. H2O2.—If an acid solution of a vanadate is treated aie a 
few drops of H2O2 and shaken, the solution becomes colored reddish 
brown owing to the formation of pervanadic acid, HVOs, insoluble 1 In F 
ether. This is a very delicate reaction. 

8. Mercurous Nitrate precipitates white mercurous vanadate from 4 
neutral solutions of a vanadate; the precipitate is soluble in nitric acid. — 

9. Oxidizing Agents convert divanadyl compounds into vanadic — 
acid. The oxidation may be effected by bromine in hydrochloric — ; 
acid solution, 


[V20o]**-+Bro-+6H20 — 2V0."-+2Br-+ 12H", . 


by dilute potassium permanganate in hot, very dilute sulfuric acid : . 
solution, 


++ | 
5[V202]* t+2Mn0.-+22H20 > 10VO4= +2Mn*+++44Ht, 


or by sodium peroxide in alkaline solution: 


++ 
[V202]* ++ Na202+8 OH- > 2VO04=+2Nat+4H20. 


Chromium, aluminium, vanadium and uranium may be separated from iron, 
nickel, cobalt and manganese by means of this reaction; the chromate is left — 
in solution as sodium chromate, NazCrO,, the aluminium as sodium aluminate, 
NaAlO,, the vanadium as sodium vanadate, Na;VO,, the uranium as sodium — 
peruranate, and the zinc as sodium zincate, Na2:ZnO,; while the iron is pre- — 
cipitated as Fe(OH);, the nickel as Ni(OH). or Ni(OH);, the cobalt as Co(OH)s, — 
and the manganese as hydrated MnO,. <j 


Detection of Vanadium in Rocks (Hillebrand) * q 


Fuse 5 gms. of the finely powdered rock with 20 gms. of Na,CO; and 3 gms. — 
of NaNO;. Cool, extract the fused mass with water, reduce the manganate — 
formed by the addition of a little alcohol, and filter the solution which contains _ 


* Amer. J. Science, 1898, p. 209. 








METALS OF THE HS GROUP 491 


the sodium salts of arsenic, phosphoric, molybdic, chromic, vanadic and tungstic 
acids. Nearly neutralize it with nitric acid (the amount necessary having been 
determined by a blank test), evaporate nearly to dryness, take up in water, 
and filter. Treat the alkaline solution with mercurous nitrate, whereby mer- 
curous phosphate, arsenate, chromate, molybdate, and tungstate with some 
basic mercurous carbonate are precipitated. Boil the solution, filter, dry the 
precipitate, separate it from the filter, ignite it in a platinum crucible, and 
fuse with a little sodium carbonate. Extract the fused mass with water, when 
a yellow color shows that chromium is present. Acidify the solution with 
sulfuric acid, and precipitate traces of Pt, Mo, and As by means of HS (best 
in a small suction flask). Filter off this precipitate and remove the excess of 
H.S from the filtrate by boiling, while passing a stream of carbonic acid gas 
through it. Evaporate the solution to dryness, and carefully expel the excess 
of sulfuric acid, heating in an air-bath. Dissolve the residue in 2 or 3 cc. of 
water and shake with a few drops of H.O.; a brownish-yellow color shows 
the presence of vanadium. If chromium is present, on adding H.O, and ether 
to the sulfuric acid solution and shaking, the ether will be colored blue by 
chromium and the aqueous solution yellow by vanadium.* 


REACTIONS IN THE DRY WAY 


The borax bead is colorless in the oxidizing flame if slightly saturated 
with the vanadium compound, yellow if strongly saturated, and green 
in the reducing flame. 


MOLYBDENUM, Mo. At. Wt. 96.0. M. Pt.=2500? 


Occurrence—Molybdenite, MoSe; wulfenite, PbMoQ.; powelite, 
CaMoO4. Molybdenum has a valence of 2, 3, 4, and 6, and forms the 
following oxides: MoO, Mo203, MoOe2, and MoOs. The first three 
are basic anhydrides, while the last oxide, MoOs, is an acid anhydride, 
forming a white mass (yellow when warm) which is readily fusible, 
but very difficultly volatile. When heated strongly, colorless, trans- 
parent, thin, orthorhombic plates of MoO3 may be obtained from 
the fumes. MoQOgz is only very slightly soluble in water, but dissolves 
readily in alkalies and in ammonia, forming molybdates. Molybdic 
acid itself can be readily obtained as a solid mass by acidifying the 
solution of an alkali molybdate; it is soluble in an excess of the acid 
(difference from tungstic acid). . The most important commercial 
molybdate is the acid ammonium molybdate, corresponding to the 
formula: 

(NH4) 6Mo7O24 +4H20. 


*E. Cuampaane, Chem. Zentralbl., 1904, II, p.-1167. 





492 REACTIONS OF SOME OF THE RARER METALS 


REACTIONS IN THE WET WAY 


A solution of ammonium molybdate should be used. 
The alkali molybdates are soluble in water; the remaining salts are mostly ' 
insoluble in water but soluble in acids. 


1. Dilute Acids precipitate from concentrated alkali molybdate — 
solutions white HzMoO,, soluble in an excess of acid. j 
Concentrated Sulfuric Acid.—If a trace of a molybdenum com- ~ 
pound is evaporated with a drop of concentrated sulfuric acid almost — 
to dryness in a porcelain dish, the mass is colored intensely blue. This — 
is an exceedingly delicate reaction. 
2. H2S at first colors acid molybdenum solutions blue, and pre- 7 
cipitates, little by little, the molybdenum as brown molybdenum tri- — 
sulfide, MoSs, soluble in ammonium sulfide, forming a brown solution — 
from which MoS83 ‘is reprecipitated by the addition of acids. Molyb- ~ 
denum sulfide is oxidized by treatment with concentrated nitric acid, — 
or by roasting in the air, into MoOs. ¥ 
3. Zinc.—If a molybdate solution which is acid with hydrochloric — 
or sulfuric acid is treated with zinc, the solution is colored at first blue, — 
then green, and finally brown. Other reducing agents such as SnCle, 
Hge(NOs3)2, etc., cause the same reaction. 
4. SOs does not reduce dilute, strongly acid solutions of molyb- — 
dates either in the cold or on heating. Neutral or slightly acid 
solutions are reduced and colored blue. “@ 
5. KCNS causes no change when added to acid molybdenum — 
solutions, but if the solution is then treated with zine or stannous 
chloride, a blood-red coloration is produced on account of the forma- 
tion of molybdenum thiocyanate; the reaction also takes place in the — 
presence of phosphoric acid (difference from iron). If the solution — 
is shaken with ether, the colored compound is dissolved in the latter. © 
6. Sodium Phosphate.—If a few drops of a solution of sodium ~ 
phosphate are added to a molybdate solution strongly acid with nitric 
acid, a yellow crystalline precipitate of ammonium phosphomolybdate \ 
is formed, slowly in the cold, but much more quickly on warming the 
solution (cf. Phosphoric Acid, p. 379). Arsenic acid causes the pre- 4 
cipitation of a similar compound (ef. p. 231). 
7. Mercurous Nitrate precipitates white mercurous molybdate E 
from neutral solutions; the precipitate is soluble in nitric acid. 2 
8. Lead Acetate pile, eg white lead molybdate, soluble in 4 
nitric acid. a 
9. Potassium viecsteahte produces a reddish brown precipitate 4 
A very sensitive test. 4 


A a 


. 





METALS OF THE H,S GROUP 493 


Molybdenum solutions containing free oxalic, acetic or phosphoric acids 
usually give no precipitate with potassium ferrocyanide, but merely a brown 
coloration. Molybdenum ferrocyanide is, however, insoluble in dilute mineral 
acids but is dissolved by concentrated hydrochloric acid and reprecipitated 
upon diluting. It is readily soluble in caustic alkali and ammonia solutions, 
in which respect it is different from the uranyl and cupric ferrocyanides (pp. 156 
and 220). To detect the ferrocyanide ion in molybdenum ferrocyanide, 
dissolve the salt in ammonia, saturate the ammoniacal solution with H,S, 
acidify with dilute H.SO,, filter off the MoS; and test the filtrate with ferric 
chloride solution (p. 150). 

10. H.O;. If a solution to be tested for molybdenum is evaporated to 
dryness on the water-bath, the residue treated with a little concentrated 
ammonia and then with hydrogen peroxide, the ammoniacal solution is immedi- 
ately turned- pink or red. Then, evaporating to dryness again and treating 
the residue with sulfuric or nitric acid, yellow permolybdic acid, HMoOQ,, 
is obtained. 


REACTIONS IN THE DRY WAY 


Alkali molybdates, alone or with sodium carbonate, are reduced 
on charcoal to gray molybdenum, a white incrustation of MoOg3 being 
formed at the same time. 

Salt of Phosphorus Bead.—All molybdenum compounds color the 
bead, but the color depends upon the concentration. In the oxidizing 
flame the hot bead is colored brownish-yellow to yellow; it becomes 
yellowish-green on cooling and finally colorless... In the reducing flame 
the bead becomes dark brown when hot and grass green when cold. 
The borax bead is similar but not quite as characteristic. 


Separation of Molybdenum, Arsenic, Antimony, and Tin 


These elements are all precipitated as sulfides upon the introduction of 
hydrogen sulfide into an acid solution. They are separated from the members 
of the copper group by treatment with ammonium polysulfide solution, in which 
their sulfides are soluble. On acidifying this solution of the thiosalts with 
dilute hydrochloric acid, the molybdenum, arsenic, antimony and tin are repre- 
cipitated as sulfides. 

Filter off this precipitate, wash, dry and introduce it, little by little, into a 
nickel crucible containing a molten mixture of 10 parts NaOs, and 10 parts 
Na,CO; for each part of sulfide precipitate. Fuse the contents of the crucible, 
after all the sulfide has been introduced, for ten minutes over the Bunsen 
burner, then cool and extract with cold water. The aqueous solution thus 
obtained may contain sodium arseniate and sodium molybdate and the insoluble 
residue may consist of sodium antimonate and tin dioxide. Filter off this 
residue and wash it with a normal solution of sodium hydroxide. Test the 
filtrate for arsenic by acidifying with hydrochloric acid, making strongly 
ammoniacal and adding magnesium mixture. A white crystalline precipitate 
of magnesium ammonium arseniate is formed if arsenic is present, but only 


494 REACTIONS OF SOME OF THE RARER METALS 


after standing for some time with a little arsenic. Saturate the filtrate from 


the magnesium ammonium arseniate, precipitate with hydrogen sulfide to con- 
vert any molybdenum present into ammonium, thiomolybdate, and then — 


acidify the solution with dilute hydrochloric acid. Filter off the precipitated 
sulfide, treat it with concentrated nitric acid in a porcelain crucible and test 
for molybdenum with concentrated sulfuric acid, as described on p. 492. 

Test for Antimony and Tin.—Treat the residue insoluble in dilute caustic 
soda solution with a mixture of equal parts concentrated hydrochloric acid and 
water, place the solution in contact with a piece of platinum foil and test for 
antimony and tin as described on p. 275. 


TUNGSTEN, W. At. Wt. 184. M. Pt.=3000° 


Occurrence.—Tungsten is not very often found in nature, but there 
are a number of well-crystallizing tungsten minerals, such as the 
minerals of the Scheelite group. 

Scheelite, CaWO.; cuproscheelite, (CaCu)WO4; reinite, FeWO,; 
stolzite, PoWO.4. These minerals all crystallize in the tetragonal system 
and form with powellite, CaMoO., and wulfenite, PbMoO., a very 
interesting isomorphous group. Another isomorphous group, which 


consists of minerals crystallizing in the monoclinic system, is formed — 


by hiibnerite, MnW0O.,; wolframite, (MnFe)WO.a, and _ferberite, 


FeWO,4. The most important tungsten mineral is wolframite, which 


is usually contaminated with small amounts of silicic, tantalic, and 
niobic acids. Tungsten forms two oxides, WO2 and WOQOs. 

WOz is a brown powder, readily obtained by heating WOz3 to dull 
redness in a stream of hydrogen. It is pyrophoric and must, therefore, 
be cooled in a stream of hydrogen before it is allowed to come into 
contact with the air. By igniting strongly in a stream of hydrogen, 
metallic tungsten is obtained, which is stable in the air. This behavior 
is important and is taken advantage of in the quantitative deter 
tion of tungsten. 

WOs is an acid anhydride obtained by the ignition of tungstic acid 
of ammonium or mercurous tungstates, or by the oxidation of the 
dioxide on heating in the air. 

The trioxide is a canary-yellow powder, insoluble in water and 
dilute acids, and only slightly soluble in concentrated hydrochloric and 


hydrofluoric acids. It dissolves readily by warming with potassium 


or sodium hydroxides, and less readily in ammonia. It is most easily 
dissolved by fusing with sodium carbonate, sodium tungstate being 
formed: 


W03+ Na2CO3 = NazW04+COrz. 





ee ee ee ee eee 


ev 2 a 


METALS OF THE H.S GROUP * 495 


It is changed to potassium tungstate by fusing with potassium 
pyrosulfate: 
W0O3+ K28207 = K2W04+ 2803. 


If the product of this last fusion is treated with water, usually none 
of the tungsten goes into solution, because if an excess of potassium 
pyrsoulfate is present (which is usually the case) it reacts with the 
potassium tungstate, forming free tungstic acid: 


K2eW04+ Ke8207+ H20 = 2K2804+HeWOg. 


If not enough pyrosulfate remains to complete the above decom- 
position, some of the tungsten will be dissolved, but never all of it. 
If a little sulfuric acid is added to the water, none of the tungsten 
will go into solution. This property enables one to separate tungsten 
from titanium. If ammonium carbonate is added, all of the tungsten 
dissolves, which enables us to separate tungstic from silicic acid. 


REACTIONS IN THE WET WAY 


A solution of sodium tungstate should be used. 

1. Mineral Acids, HCl, HNOz, H2SO.4, produce, in the cold, a 
white, amorphous precipitate of hydrated tungstic acid, H2W0O4+ H20.* 
By boiling the solution, the yellow anhydrous acid HzWOs is obtained, 
insoluble in dilute acids, but soluble to an appreciable extent in con- 
centrated hydrochloric acid. 

Tungstic acid must always be washed with water which contains 
acid or a dissolved salt, as otherwise tungstic acid will form a pseudo- 
solution with pure water, so that a turbid filtrate will be obtained 
(ef. pp. 58, 127 and 218). 

Phosphoric acid behaves differently toward solutions of the alkali 
tungstates than do the other mineral acids; it produces a white pre- 
cipitate soluble in an excess of phosphoric acid; a complex phospho- 
tungstic acid is formed, e.g., NagPO4-12WO3. If the solution of an 
alkali tungstate is boiled with free tungstic acid, the latter gradually 
goes into solution, forming a metatungstate; 


NaeWO,4 +3WOz3 = NaeW.40i13. 


Mineral acids cause no precipitation in solutions of metatung- 
states. If the solution is boiled with an excess of acid, the soluble 
metatungstic acid is gradually changed to insoluble, ordinary tungstic 
acid, which is then precipitated. 





* The presence of tartaric acid prevents the precipitation. 


496 REACTIONS OF SOME OF THE RARER METALS 


2. HeS produces no precipitation in acid solutions. . 
3. (NH4)2S gives no precipitation in a solution of an alkali tung- | 
state, but if the solution is afterward acidified, light-brown tungsten — 
trisulfide, WS3, is precipitated, which has the property of forming — 
pseudo-solutions with pure water, but is insoluble in hydrochloric acid. : 
The precipitate redissolves in ammonium sulfide. ; 
4. Reducing Agents.—If the solution of an alkali fuga is 
treated with HCl and zinc, the tungstic acid at first precipitated by 
the HCl is soon turned to a beautiful blue color, owing to the forma- 
tion of W20s. , 
SnCle produces a yellow coloration at first, but on adding HCl . 
and warming, a beautiful blue precipitate is bistied: This is one of - 
the most sensitive reactions for tungstic acid. : a 
5. Mercurous Nitrate precipitates white mercurous tungstate — 
from neutral solutions. 
6. Lead Acetate precipitates .white lead tungstate from neutral — 
sclutions. 4 


REACTIONS IN THE DRY WAY 


The salt of phosphorus bead is colorless in the oxidizing flame, 
and blue in the reducing flame, becoming blood red on the addition 
of alittle FeSO. | 


SELENIUM, Se. At. Wt. 79.2. Sp. Gr.=4.28-4.5. M. Pt.=217-220° C, 


_Occurrence.—Although selenium is quite widely distributed in nature, — 
it is invariably found in very small amounts, usually replacing sulfur, — 
forming isomorphous compounds with lead, silver, copper, and mercury; — 
clausthalite, PbSe; berzelianite, (CuAgTl)2Se; naumannite, (Age Pb)Se; — 
tiemannite, HgSe; lehrbachite, (Pb,Hg)Se; onofrite, Hg(SeS); eu- — 
eairite, (Ag,Cu)2Se. It is also found in small amounts in many varieties — 
of pyrite and chalcopyrite, and indeed the small quantities which are — 
found in these minerals form the chief source of the selenium of com- — 
merce. By roasting these minerals (as in the manufacture of sulfuric — 
acid) all of the selenium is volatilized, and is consequently deposited — 
in the lead chambers as a mud from which it is extracted with a solution — 
of potassium cyanide and afterwards precipitated with acid: 


KCN+Se=KCNSe and KCNSe+HCl=HCN+KCI-+Se. 


Selenium, like sulfur, exists in two allotropic forms. The modi- 4 
fication soluble in carbon disulfide is obtained by reducing selenious acid — 
in the cold with sulfurous acid; it is a brick-red powder. After heating © 





METALS OF THE H,S GROUP 497 


this red selenium with hot water for some time, it is changed into black 
selenium, and is then insoluble in carbon disulfide. 

On heating in the air, selenium burns. with a bluish flame (giving 
off an odor similar to that of rotten radishes) forming white, crystalline 
selenium dioxide, SeO2, which will sublime on being heated in a stream 
of oxygen. . Selenium forms one oxide, SeOz, and two acids: selenious 
acid, H2SeOs, and selenic acid, H2SeOx. 

Selenious acid, H2SeO3, is obtained in the form of long colorless 
needles by oxidizing selenium with nitric acid or aqua regia,* or by 
the solution of its anhydride, SeOe, in water. Unlike sulfurous acid, 
it is not changed on standing in the air into selenic acid; but, on the 
contrary, is reduced by dust, etc., to red selenium. The acid is dibasic, 
and forms salts in which either one or both of the hydrogen atoms are 
replaced by metal. 

The acid salts are all soluble in water, but the neutral salts are all 
insoluble with the exception of those of the alkalies. 

Selenic acid, H2SeOz, is obtained in solution by conducting chlorine 
into water which contains either suspended selenium or dissolved 
selenious acid: 

Se+3Cle+4H20 = H2Se04+6HCI. 


Sodium seleniate is obtained by fusing selenium with sodium car- 
bonate and potassium nitrate. Selenic acid is a dibasic acid and 
behaves similarly to a peroxide, evolving chlorine when boiled with 
concentrated hydrochloric acid, being reduced to selenious acid: 


- HeSe04+2HCl=H.0+H28e03+Cl f. 


REACTIONS IN THE WET: WAY 


(a) Selenious Acid 
A solution of either potassium selenite or of free selenious acid 
should be used. 


1. HeS produces a lemon-yellow precipitate, consisting of selenium 
and sulfur, from solutions in water or in dilute hydrochloric acid: 


H2SeO03 + 2H2S = 3H2O +Se +28. 


The precipitate is soluble in ammonium sulfide. 
2. Reducing Agents. 
SOz2 precipitates red selenium. The solution may be hot or cold, 





* If a solution of selenium in aqua regia is evaporated, considerable selenium 
is lost by volatilization; the addition of KCl or NaCl to the solution prevents 
such loss. 


- 


498 REACTIONS OF SOME OF THE RARER METALS ~ 





and contain little or much hydrochloric or sulfuric acid; by long boiling q ; 
the precipitate turns grayish black. a 
SnCly precipitates red selenium even in the presence of considell q 
able sulfuric acid. 3 
FeSO. immediately precipitates selenium from concentrated solu- — 
tions of selenious acid containing hydrochloric acid; from dilute solu- — 
tions the precipitation takes place very slowly and incompletely if 
much sulfuric acid is present. | 
Hydroxylamine hydrochloride precipitates selenium from solu- — 
tions of selenious acid containing sulfuric or hydrochloric acid on long — 
boiling; the precipitated selenium is red at first but finally becomes — 
gray (difference from tellurium). 4 
Hydrazine hydrochloride precipitates selenium from hot acid — 
and alkaline solutions; the selenium is red at first and finally gray. 
Hydriodic Acid (KI and HCl) Dt ag a red selenium in the 4 
cold (difference from tellurium). % 
Zinc precipitates red selenium from acid solutions: the zine becomes — 
coated with Se and looks as if covered with Cu. 4 
3. BaCle precipitates from neutral solutions white barium selenite, 
BaSeOs, soluble in dilute acids. a 
4. CuSO. produces a greenish-blue, crystalline precipitate (differ- 4 
ence from selenic acid). 


(b) Selenic Acid 


1. HeS causes no precipitation unless the solution is boiled with 
hydrochloric acid. In the latter case the selenic acid is reduced first to 
selenious acid and then to selenium, which precipitates together with . 
free sulfur. 

2. BaCl2 gives a white precipitate of barium seleniate, BaSeQOu, 
insoluble in water and in dilute acids, soluble, with evolution of chlorine, 


| 
A solution of potassium seleniate should be used. 
on being boiled with hydrochloric acid: 


BaSe0.4+4HCl=BaCle+ H2Se03+Cle fT +H20. 


3. CuSO, produces no precipitation. 
4. SO does not reduce selenic acid except by long boiling with — 
hydrochloric acid. The reduction takes places more readily with — 
hydrazine. 





METALS OF THE HS GROUP 499 


Method for Testing Sulfuric Acid for Selenium * 


Add 5 or 6 drops of the acid to be tested to a freshly prepared solution 
of a little codein in sulfuric acid; if selenium is present, a green coloration will 
be apparent. The test is a very delicate one. 


REACTIONS IN THE DRY WAY 


All selenium compounds emit the odor of decayed radishes on being 
mixed with sodium carbonate and heated on charcoal before the blow- 
pipe. 

If a selenium compound is heated at the end of a thread of asbes- 
tos in the upper reducing flame of the Bunsen burner, it will be reduced 
to selenium; and if a test-tube filled with water is held above the flame, 
a red coating of selenium will be deposited upon the glass.t If a few 
drops of concentrated sulfuric acid { are placed in a larger test-tube 
(large enough to hold the smaller test-tube) and the tube on which the 
selenium is deposited is emptied and placed within the’ larger tube, 
the selenium will dissolve § in the sulfuric acid, forming a green solu- 
tion; but on the addition of water, red selenium will be reprecipitated 
(difference from tellurium) : 


SeSO3-+ H20 =Se+He2S0O.. 


Green. 


> 


TELLURIUM, Te. At. Wt. 127.5. Sp.Gr.=6.1-6.4. M. Pt.=452° C. 


Occurrence.—Tellurium is a rarer element than selenium, always 
occurring in. the form of a telluride, and usually combined with the 
noble metals: calaverite, (Au,Ag)Te2; krennerite, (Au,Ag)Te2; sylvan- 
ite, (Au,Ag)Tes4; nagyagite, AugzSbePbioTesSi5; coloradoite, HgTe; 
silver telluride, AgeTe; and often in small amounts in galena and copper 
ores. Emmonsite of Cripple Creek, Colorado, is a ferric telluride with 
‘70.71 per cent TeOo and 22.76 per cent Fe2O3. Tellurium itself is 
a bluish-white, brittle substance, which can be distilled in a stream of 
hydrogen. It burns in the air with a bluish-green flame, forming 
tellurium dioxide, TeOz. It is insoluble in carbon disulfide, and can 





. DRAGENDORIF, Chem. Zentralb., 1900, 944. 

+ Cf. p. 68. : 

t The sulfuric acid should be freed from water by heating in a platinum crucible 
to a temperature just below the boiling-point, and the crucible hin its contents 
allowed to cool in a desiccator. 

§ Slowly in the cold, readily on warming. 


500 REACTIONS OF SOME OF THE RARER METALS 


be oxidized by means of nitric acid to tellurous acid. On being fused 
with potassium cyanide, out of contact with the air, it is changed to 
potassium telluride, 


2KCN +Te = KoTe+(CN)s2, 


which dissolves in water, forming a cherry-red solution. If air is con- 
ducted through this solution, the tellurium is. precipitated in the form 
of a black powder (difference from selenium) : 


KeTe+H20+0=2KOH+Te. 


Tellurium may be separated from selenium by means of this last 
reaction. The two metals are fused with potassium cyanide, the melt 
is treated with water, and the tellurium precipitated by passing a current 
of air through the solution; the selenium is precipitated from the filtrate 
by acidifying with hydrochloric acid. Tellurium forms two oxides: 
TeOz and TeQs3. 


Tellurium dioxide (the anhydride of tellurous acid) is usually obtained in 
the form of a white mass, which melts on gentle heating, forming a yellow 
liquid. Tellurium dioxide does not sublime (difference from selenium). It is 
scarcely soluble at all in water, is slightly soluble in ammonia and in dilute 
acids, but readily soluble in concentrated acids or in caustic potash solutions. 
TeO, dissolves in fairly concentrated sulfuric acid, forming the basic sulfate, 
Te.0;-SO., while with nitric acid it forms the basic nitrate, TexO;(OH)NO;. 
Both of these compounds are hydrolized readily, forming insoluble tellurous 
acid; the latter in turn loses water and forms the anhydride. 

On dissolving TeO. in caustic potash, potassium tellurite, K,TeO;, is 
obtained. Only the alkali tellurites are soluble in water. 

Tellurium trioxide (telluric anhydride) is formed by heating tellurie acid. 
It is a yellow powder, insoluble in water and nitric acids, scarcely affected by 
boiling with concentrated hydrochloric acid, but is readily dissolved by boiling 
with a concentrated solution..of potassium hydroxide (but not by sodium 
hydroxide), forming potassium tellurate. 

Telluric acid, H,.TeO.,+2H;0, is a very weak acid, obtained by oxidizing 
tellurous acid with chromic acid, and precipitating the telluric acid by the 
addition of concentrated nitric acid. The acid forms a colorless crystalline 
mass, is readily soluble in water, and is converted by means of concentrated 
hydrochloric acid into tellurous acid, with evolution of chlorine. The acid 
dissolves readily in caustic potash (or soda) solution, forming the readily 
soluble alkali tellurate, which reacts strongly alkaline in aqueous solution. 

By gently heating the hydrated telluric acid, the anhydrous acid, H:TeQ,, 
is obtained in the form of a white powder and is totally different from the 


hydrated acid. The latter is soluble in water and in caustic alkalies, and is ~ 
completely reduced by boiling with concentrated hydrochloric acid; but the — 


anhydrous acid is insoluble in water and in concentrated sodium hydroxide 
solution, and is only very slightly attacked by boiling, concentrated hydro- 


vr ee 


_—_— wees Se 5 hattatlmn, — = 


7 , oe 


Es ee ee a are eer nS) aS ee ee ee A 


ES a 








METALS OF THE H,S GROUP 501 


chloric acid, although readily soluble in warm potassium hydroxide solu- 
tion. 

Only the alkali tellurates are soluble in water; the others are usually 
obtained in the form of amorphous precipitates soluble in acids. 


REACTIONS IN THE WET WAY 
(a) Tellurous Acid 
A solution of potassium tellurite, K,TeO;, should be used. 


1. HeS precipitates from acid solutions brown TeS2, which is readily 
soluble in ammonium sulfide. 

2. Reducing Agents. 

SO» precipitates tellurium completely from dilute hydrochloric 
acid solutions in the form of a black powder, even in the presence of 
tartaric acid; but from a solution containing considerable hydro- 
chloric acid no tellurium is precipitated even opr boiling (difference 
from selenium). The separation of the selenium from tellurium 
can be accomplished in hydrochloric acid, sp. gr. 1.18. 

SnCl2 or Zn causes black tellurium to precipitate from solutions 
tvhich are not too acid. 

H3PO3 precipitates the tellurium only from concentrated solu- 
tions, not at all from cold dilute solutions. 

FeSO, reduces neither tellurous nor telluric acids (difference from 
selenium). 

Hydroxylamine Hydrochloride produces no precipitate in solutions 
of tellurous acid containing mineral acids, but precipitates tellurium 
completely by boiling the ammoniacal solutions for a long time. 


INH.OH+Te037+2H*t > 4H20+N20 T +Te. 


Hydrazine Hydrochloride precipitates black tellurium both from 
acid and ammoniacal solutions: 


NeoH4:-2HCI+TeO3 — 3H20+ Ne T +Te+2CI. 


Hydriodic Acid (KI and HCl) produces no precipitation, but merely 
a reddish-brown coloration that turns light yellow on boiling (differ- 
ence from selenium). 

Zinc (or Fe, Sb, Sn, Cd, Hg, Pb, Cu, etc.) COANE black tellu- 


* yium. 


3. HCl produces a white euscinabata of HeTeQ3. 


502 REACTIONS OF SOME OF THE RARER METALS 


(b) Telluric Acid. 


A solution of potassium tellurate should be used. 

1. HCl causes no precipitation; but if the solution is bailed chlorine 
is evolved, and on dilution with water tellurous acid is precipitated. | 

2. HeS and reducing agents have the same effect upon hot solu- 
tions of tellurates as upon tellurites. 

3. Lead Salts precipitate difficultly soluble lead tellurate. 

4. Hydrazine Hydrochloride precipitates all the tellurium as a 
black powder by long boiling of the acid or alkaline solution: 


3(N2H4:-2HCl) +2Te04" — 8H20+2H*+6Cl-+3Ne2 Tf +2Te. 


Detection of Selenium and Tellurium in Ores 


Principle-—The finely powdered, dry ore is heated in a current of chlorine; 
the chlorides of sulfur, selenium, arsenic, antimony and iron are volatilized 
and may be absorbed in dilute hydrochloric acid. 

Procedure.—Place the finely powdered, dry mineral in a porcelain boat, 
push this into a tube of difficultly fusible glass and connect the tube on one 
side with the wash-bottle and on the other side with a 10-bulb Meyer tube. 

Prepare chlorine gas from a Kipp generator containing chloride of lime 
and hydrochloric acid, wash the gas by passing it through a bottle contain- 
ing water and dry it by passing it through concentrated sulfuric acid. Pass 
the chlorine gas through the tube and when all the air has been expelled, 
begin heating the substance, at first very gently. Fumes quickly begin to 
form, showing that the action of chlorine upon the substance has started. 
Soon vapors of sulfur chloride, 8.Cl, (B. P. 64°) begin to condense in the 
front end of the tube in the form of drops; drive these over into the receiver 
by carefully heating the tube. As the temperature is raised a little, a white 
sublimate of selenium chloride, SeCl, (B. P. 200°) forms in the front end of 
the tube; drive this over into the receiver in the same way. Now heat the 
boat hotter and soon brown vapors of ferric chloride will be evolved, some of 
these will condense, forming glistening scales which must also be driven over 
into the receiver. Continue heating until finally no more vapors are evolved. 
Transfer the contents of the 10-bulb tube to a porcelain dish, add 0.5 gm. of 
potassium chloride to prevent loss of selenium chloride by volatilization, and 
evaporate the solution to dryness on the water-bath. Dissolve: the residue in 
as little hydrochloric acid as possible and treat the solution with stannous 
chloride. If tellurium is present a black precipitate is obtained which may 
also contain selenium. Filter through an asbestos filter, wash the residue with 
dilute hydrochloric acid, place the filter and asbestos in a small test-tube, boil 
it with concentrated hydrochloric acid until no more black spots are visible in 
the asbestos, dilute with water and filter. LEvaporate this filtrate to dryness 
on the water-bath, dissolve the residue in 10 ce. of HCl, sp. gr. 1.175, and pass — 
SO, gas into the hot solution. Filter off any precipitate of red selenium that 
may form, Dilute the filtrate with considerable water and again pass SO, into 





METALS OF THE H,S GROUP 503 


the hot solution, which should precipitate the tellurium as a black powder. 
Identify the selenium and the tellurium by the sail reactions given on pp. 499 
and below. 


REACTIONS IN THE DRY WAY 


Metallic tellurium is formed by heating any telluride in the upper 
reducing flame, and can be collected on the lower surface of a test-tube, 
which is filled with water, in the form of a black film, soluble in 
concentrated sulfuric acid. The latter solution is of a carmine-red 
color (difference from selenium); on the addition of water black 
tellurium is deposited: 


TeSO3 +H:0O =Te -|- He2SO.. 


Carmine red 


THE PLATINUM METALS 


PLATINUM, PALLADIUM, RHODIUM, OSMIUM, RUTHENIUM AND 4 
IRIDIUM ; 


Platinum has been described a'ready on p. 266. 


PALLADIUM, Pd. At. Wt. 106.7. Sp. Gr.=11.8. M. Pt.=1550° C. 


Occurrence.—The platinum metals form an isodimorphous group, — | 


but only in the case of palladium are both forms known—the isometric 
and the hexagonal: 


(a) Isometric System. (b) Hexagonal System. 
Platinum (Pt, Fe). Tridosmium (Sysserskit) (Ir, Os). % 
Iridium (Ir, Pt). Osmiridium (Newjanskit) (Ir, Os, 
Platinum iridium (Pt, Ir, Rh). Pt, Rh, Ru) or (Os, Ir, Rh). 
Palladium (Pd, Pt, Ir). Palladium (Pd, Pt, Ir). 


Properties—Rolled, hammered, or cast palladium possesses an 
almost silver-white color, but when precipitated from solutions it is 
in the form of a black powder. If it is suspended in water when in 
the finely divided form, it is transparent with a reddish color. Palla- 
dium has the lowest melting-point of all the platinum metals, viz., 
1550° C. On being heated in the air, it appears bluish, owing to the 
formation of PdeO; the latter, however, is decomposed by stronger — 
heating. 

Behavior towards acids: Although the other platinum metals are 
attacked by no acid except aqua regia, palladium is dissolved slowly 


by warm nitric acid (also in the cold when it is alloyed with other — 


metals such as Cu, Ag, etc.), forming a brown solution of Pd(NQs3)o. 

Finely divided, precipitated palladium is soluble in hydrochloric 
acid wher exposed to the action of air at the same time, and less readily 
soluble in sulfuric acid. It is readily attacked by fusing with potas- 
sium pyrosulfate, forming soluble palladium sulfate, PdSOz. — 

The best solvent for palladium is aqua regia. 

Finely divided palladium has the very characteristic property of 
being able to absorb almost 700 times its own volume of hydrogen, 
and possesses consequently a very strong catalytic action. If hydro- 

504 





THE PLATINUM METALS 505 


gen and oxygen (air) are conducted at the same time over some gently 
ignited, finely divided, metallic palladium, the hydrogen is burnt to 
water, and in the same way carbon monoxide may be changed to carbon 
dioxide. Methane, however, is only decomposed by igniting the 
palladium more strongly, so that this gives us a method for separa- 
ting methane from a mixture of H and CO (cf. Vol. 2, Gas Analysis). 

Colloidal palladium preparations show marked catalytic effects. 
Thus by passing hydrogen gas into solutions of unsaturated crganic 
compounds a direct reduction (hydrogenation) often takes place if 
a little colloidal palladium is present. 

Palladium forms two oxides, both of which posses$ strongly basic 
properties: PdO and PdOzg. From the former the palladous, and from 
the latter the palladic, compounds are derived. The palladous com- 
pounds are much more stable than the palladic compounds, and the 
latter constantly exhibit the tendency to change into the former. 

By dissolving finely divided palladium in hydrochloric acid, palla- 
dous chloride is formed; or, better, by dissolving the metal in aqua 
regia, in which case a mixture of palladous and palladic chlorides is 
at first obtained. If this solution, however, is evaporated to dryness, 
palladic chloride loses chlorine and is completely changed into palla- 
dous chloride, so that on treating the residue with water a solution of 
palladous chloride is obtained. Since palladic chloride is decomposed 
completely by evaporation, it is bday that palladic chloride cannot 
exist in hot solutions. 


REACTIONS IN THE WET WAY 


(a) Palladous Compounds 
Use a solution of palladous chloride, PdCh. 


1. H2S precipitates black palladous sulfide from acid and neutral 
solutions. The precipitate, PdS, is insoluble in ammonium sulfide, 
but soluble in boiling hydrochloric acid, or more readily in aqua regia. 

2. KOH or NaOH precipitates a brown basic salt, soluble in an 
excess of the reagent. If the solution is acidified with HCl, then 
KOH produces no precipitate (difference from platinum). 

3. NaCOz produces a brown precipitate of palladous hydroxide, 
Pd(OH)s, soluble in excess but reprecipitated on boiling. 

4. NH4OH gives a flesh-colored precipitate of [Pd(NH3)2Cle], * 
soluble in an excess of ammonia, forming a colorless solution (eon- 
taining palladodiamine chloride, Pd(NH3)4Cle), from which yellow 





* This compound is an isomer of palladosamine chloride, and is often written 
thus: PdCh,Pd (N Hs) aCle. 


506 REACTIONS OF SOME OF THE RARER METALS | 


crystalline palladosamine chloride, Pd(NH3)2Cle, is precipitated on the 
addition of hydrochloric acid. The latter compound is difficultly — a 
soluble in dilute hydrochloric acid and is used for the preparation of — 
pure palladium. pris 

In a solution of palladous nitrate, ammonia causes no precipitation, 
but forms colorless palladodiamine nitrate, Pd(NH3)4(NO3)2. 

5. NH,Cl. If a solution of palladous chloride or of sodium-palla- 
dous chloride is treated with ammonium chloride and evaporated to 
dryness on the water-bath, the residue is soluble in a very little water. 
If the solution is acidified with nitric acid, gradually all the palladium 
is precipitated as red (NH4)2[PdCle] (difference from platinum). 

6. KCl, when added to a concentrated solution, causes the pre- 


cipitation of difficultly soluble, reddish-brown Ke{[PdCl4] (octahedrons). C 


7. HI or KI produces a black precipitate of palladous iodide, even 
in very dilute solutions. The precipitate is insoluble in water, alcohol, 
ether, and HI, but soluble in KI and NH3. (This and the following 
reaction are characteristic of palladium.) . 

8. Hg(CN)2 produces a _ yellowish-white gelatinous precipitate 
of palladous cyanide, Pd(CN)e, difficultly soluble in HCl, readily 
soluble in KCN and NH3. On being ignited, the spongy metal remains. 

9. Nitroso-6-naphthol (a saturated solution in 50 per cent acetic 
acid) gives a voluminous, brown precipitate of Pd(CiopHegNO2)2 even 
in the most dilute solutions (difference from platinum). 

10. Reducing Agents.—H2SOs, formic acid, HCOOH, Zn, Fe, 
FeSO4, CueCle,* alcohol, and CO7f reduce palladium salts to the 
metal itself. 

In the presence of HCl, stannous chloride forms at first a red, then 
a brown, and finally a green, solution; but in the absence of the acid, 
SnCle causes a partial reduction to the metal and the solution turns 
green. 


(b) Palladic Compounds 


These give the same reactions as palladous compounds, on account 
of their being readily changed into the latter. The principal differ- 
ence, however, is the insolubility of the ammonium salt of chloropalladic 





* In the presence of considerable NaCl or HCl there is no reduction with CusCh. 
+ PdCl,+CO+H,0 =2HCl1+CO;+Pd. This reaction enables one to detect 
small amounts of CO in gas mixtures; e.g., in the air. For this purpose the gas 
is led through a narrow glass tube into 10 cc. of a solution which contains 1 mgm. 
of PdCl, and 2 drops of dilute HCl. If CO is present, black Pd will be deposited, 
and the solution will become decolorized little by little. (Porarin and Drourn, 


Compt. rend., 126, 938.) If too much HCl is present the reduction will not take __ 


place unless NaC,H;O, is added. 





THE PLATINUM METALS 507 


acid. If a concentrated, cold solution of palladous chloride is shaken 
with chlorine water and then treated with ammonium chloride, a red 
crystalline precipitate of (NH4)2[PdClg] is soon formed. 


REACTIONS IN THE DRY WAY 


All palladium compounds are decomposed on ignition, leaving 
behind the metal, which is soluble in nitric acid or in aqua regia, and 
_the solution thus obtained can be tested by the above reactions. 


RHODIUM, Rh. At. Wt. 103.0. Sp.Gr.=12.6. M. Pt.=1920°? 


Properties.— Rhodium possesses the color and luster of aluminium; it 
is more infusible than platinum and melts at about 1920°; on cooling 
the hot metal it sputters and appears bluish, owing to oxidation. The 
solubility of rhodium depends entirely upon the fineness of the material. 

When precipitated from a solution of its chloride by means of 
formic acid or other reducing agents at a temperature not exceeding 
100°, it exists in an extremely finely divided state (rhodium black) 
and dissolves readily in boiling, concentrated sulfuric acid, or more 
readily in aqua regia. If, however, the finely divided metal is ignited 
strongly, it becomes (like the compact metal) almost insoluble in aqua 
regia. | 

If rhodium is alloyed with large amounts of other metals (Pb, Zn, 
Bi, Cu, etc.), it is left in a finely divided condition after treatment of 
the alloy with acids, and is consequently soluble in-aqua regia. When 
it is alloyed with much platinum or palladium a considerable amount 
of rhodium will dissolve in aqua regia; but when it is alloyed with a 
little platinum, most of the rhodium and a part of the platinum remain 
undissolved. 

On being fused with potassium pyrosulfate, potassium rhodium 
sulfate is formed, which dissolves in water, forming a yellow solution, 
but becomes red on the addition of HCl. : 

Rhodium forms three oxides: RhO, RhzO3, and RhOg; all possess 
a well-defined basic nature. The sesquioxide, Rh2O3, alone * forms a 
series of salts, of which sodium-rhodium chloride is the most important 
for the analytical chemist; when in this form it is easiest to bring 
rhodium into solution. This salt is prepared by mixing the finely 
divided metal very intimately with twice as much dry sodium chloride, 





* A sodium-rhodium sulfite of the formula 4Rh(SO3),6Na:sS0;+9H20 was 
prepared by Bunsen. 


508 REACTIONS OF SOME OF THE RARER METALS 


placing it in a porcelain boat and gently igniting it in a current of 
moist chlorine gas. The salt thus formed has the composition 
Nag{RhCl¢] and is soluble in water (45 parts of water dissolve 1 part 
of the salt). From this solution large, dark-red, glistening triclinic 
prisms of Na3[RhCle6]+9H20 can be crystallized out. 


REACTIONS IN THE WET WAY 
Use a solution of sodium rhodium chloride Na;[RhClk.] 


1. H2S precipitates (very slowly in the cold, but much more quickly- 
on warming) black rhodium sulfide, Rh2S3; insoluble in (NH4)28, 
soluble in nitric acid. 

2. KOH and NaOH produce at first no precipitate; but after — 
standing some time a yellow precipitate of rhodium hydroxide, 
Rh(OH)3+H20, separates out. The precipitate is soluble in an excess 
of the reagent, but it is reprecipitated on boiling in the form of brown- 
ish-black Rh(OH)3. 

In a solution of potassium rhodium sulfate, KOH precipitates the | 
yellow compound immediately. 

On adding KOH to a solution of rhodium chloride, at first no 
precipitate is produced; but on the addition of a little alcohol brown- 
ish-black rhodium hydroxide is deposited. 

3. NH4OH produces (in concentrated solutions and after stand- 
ing some time) a yellow precipitate of chlorpurpureorhodium chloride, 
Rh(NHs3)5Cls, insoluble in hydrochloric acid. 

4. KNOs, on being warmed with sodium rhodium chloride solu- 
tion, causes the precipitation of difficultly soluble, orange-yellow 
Ks[Rh(NO2)¢6], soluble in HCl. 

5. Reducing Agents.—Formic acid in the presence of ammonium 
acetate, precipitates the black metal, as does zinc in the presence of 
acids. 


REACTIONS IN THE DRY WAY 


All rhodium compounds are reduced to metal on being heated in 
a stream of hydrogen, or by heating on charcoal with sodium car- 
bonate before the blowpipe. The metal is easily recognized by its 
insolubility in aqua regia, its being brought into solution by fusing 
with potassium pyrosulfate and then treating with water, and by the 
formation of the brown hydroxide when.KOH and a little alcohol are 
added to the solution thus obtained. 





THE PLATINUM METALS 509 


- 


OSMIUM, Os. At. Wt. 190.9. Sp. Gr.=21.3-22.48. M. Pt.=about 
2700° 


Osmium and ruthenium are distinguished from the other platinum 
metals by their forming volatile oxides. 

Properties—The compact metal possesses a bluish-white color, very 
similar to zine, and is the heaviest of all metals. It can be melted by 
heating in an electric furnace.* Very finely divided osmium is oxidized 
by the air + at ordinary temperatures, and at about 400° C. it ignites 
and burns rapidly to OsO4, which is volatile at 100° C. The denser 
the metal the higher the temperature necessary to effect the oxidation. 

Behavior towards Acids—In the compact condition osmium is in- 
soluble in all acids; but in the finely divided state (as obtained by 
treating its zinc alloy with nitric acid) it is soluble in nitric acid, more 
soluble in aqua regia, and most soluble in fuming nitric acid, forming 
osmium tetroxide; the latter can be separated from the solution by 
distillation. 

Compact osmium is brought into solution by fusing with NaOH and 
either KNOs or KClO3. The melt contains a salt of perosmic acid 


(OsO4). 

Osmium forms five oxides: 

OsO Os20sz, OsOg, 

Osmious oxide, grayish-black, Osmium sesquioxide, black, Osmic oxide, black-gray, 

insoluble in acids insoluble in acids insoluble in acids 

[OsOs], OsO. 
Osmic acid, known only in Perosmic acid, colorless needles 
derivatives soluble in water 


Osmium tetroxide, OsO4, (the anhydride of perosmic acid,) is the 
most important osmium compound in the eyes of the analytical chemist, 
and is obtained by the oxidation of the substance in the air by dis- 
solving the finely divided metal in fuming nitric acid or in aqua regia, 
or by fusing with NaOH and KNOgz or KCI1Os, treating the melt with 
nitric acid and distilling. Osmium tetroxide is a colorless, crystalline 
mass which sublimes at a comparatively low temperature and melts, 
forming colorless vapors at 100° C. The vapor has a chlorine-like 
odor, attacks the mucous membrane, and is poisonous. 

The chlorides of osmium can be obtained only in the dry way; 
OsCle, OsCls, OsCla are known. ‘The potassium salt of the hypothet- 
ical hydrochlorosmic acid, H2OsCle, forms dark-red octahedrons, soluble 





*F. Myuius and R. Dintz, Ber., 1898, 3187. 
7 Cf. Or. Suc, Z. anorgan. Chemie, 19, 332. 


510 REACTIONS OF SOME OF THE RARER METALS 





in water and decomposed by boiling the solution. By heating finely — a 
divided osmium with KCl in a current of chlorine, K2[OsCle] is formed; — ; 
~ it dissolves in cold water, forming a red solution. 


REACTIONS IN THE WET WAY 
Use a solution of K,[OsCl,]. 


1. If a solution of osmium chloride is treated with dilute nitric 
acid, the mixture distilled from a small retort, and the vapors received — 
in caustic soda solution, the latter will be colored yellow, owing to the 
formation of potassium osmiate. If this solution is now acidified, — 
osmium tetroxide is set free, and can be recognized by its very penetra- 
ting odor. On adding a little sodium thiosulfate to the acid solution 
and warming, a brown precipitate of osmium sulfide is formed. ‘ 

2. HeS precipitates brownish-black osmium sulfide, insoluble 
in ammonium sulfide. ay ‘ 

3. KOH, NH4sOH or K2COz precipitate reddish-brown osmium _ 
hydroxide Os(OH)a. 

4. Reducing Agents.—If the solution of the chloride is treated 
with tannic acid and alcohol and a little hydrochloric acid is added, — 
it is colored dark-blue, owing to the formation of osmium dichloride, — 
OsCle; KI colors the solution a deep reddish-purple. 

5. Indigo is decolorized by solutions containing OsO4. Ferrous 
sulfate precipitates black osmium dioxide; stannous chloride produces 
a brown precipitate soluble in HCl, forming a brown solution. 


REACTIONS IN THE DRY WAY 


All osmium compounds are reduced to metal on being heated in 
a stream of hydrogen. 


RUTHENIUM, Ru. At. Wt. 101.7: Sp. Gr.=12.26, crystallized; 11.0, 
fused. M. Pt.=above 1950° 


Properties.—Ruthenium exists in the form of a dark gray or black 
powder, and in the form of bright porous sticks; it is brittle, can be 
powdered, and is melted in the oxyhydrogen flame. 

On being melted, a part of the ruthenium is oxidized to ruthenium 
tetroxide, a volatile substance having a penetrating odor similar to 
that of OsO4. The molten metal spurts on cooling. . 

Behavior towards Acids—Ruthenium is almost completely insoluble — 
in all acids, even aqua regia. By fusing with KOH and KNOg3 (or _ 
KCIO3) it is oxidized to potassium rutheniate, K2[/RuQO,]. a 


THE PLATINUM METALS oll 


On heating with NaCl in a current of chlorine, soluble Ke[RuCle] 
is formed. The solution in water of the greenish-black melt is of an 
orange-yellow color, and colors the human skin black. Ruthenium is 
unaffected by fusing with potassium pyrosulfate. 

If forms the following oxides: 


RuO, RuzOz, RuOe, [RuOs], [RueO7z], RuOs. 


The most important of the oxides is RuO4. It is formed: 

(a) By roasting the metal itself, or its oxide, above 1000° C. (osmium 
forms the volatile tetroxide at 400° C.). 

(b) By fusing the metal with KOH and KNOs in a silver crucible, | 
dissolving the melt in water, saturating the cold solution with chlorine 
gas, and distilling the solution from a small retort: 


KeRu04+Cle = 2KCl1+RuO.. 


(c) By treating the solution of potassium-ruthenium chloride with 
KOH and Cl, and subsequently distilling. 

(d) By distilling potassium-ruthenium chloride with KClO3 and 
HCl. 

By distillling a dilute solution after the addition of nitric acid, ne RuO4 
will be evolved* (difference from osmium), 

Ruthenium tetroxide forms gold-yellow, glistening, orthorhombic 
needles that are very volatile and emit a characteristic odor; it boils 
at 100° C. and is only slightly soluble in water. It is changed by the 
addition of alcohol and HCl into ruthenium trichloride, RuCls (or 
sesquichloride, RugCle). If the solution of the latter salt is made 
ammoniacal, treated with sodium thiosulfate and warmed, an intense 
reddish-violet coloration will be produced. (This is a very sensitive 
and characteristic reaction.) 

On treating a solution of potassium ‘rutheniate with nitric acid, 
black Ru(OH)s3 is precipitated; it dissolves in hydrochloric acid, form- 
ing a yellow solution of RuCls. . 


REACTIONS IN THE WET WAY 


Use a solution of RuCl,. 


1. H2S produces no precipitation at first, but after some time the 
solution becomes azure-blue, and brown ruthenium sulfide is precipi- 
tated (very sensitive and characteristic). 





* In the presence of concentrated nitric acid, however, RuO, is formed: 


2K,Ru0O,+4HNO; =4KNO;+Ru(OH).,+ RuQ,. 


512 REACTIONS OF SOME OF THE RARER METALS 


2. (NH,)2S precipitates the brownish-black sulfide, difficultly — 


soluble in an excess of the reagent. 


3. KOH and NaOH precipitate black ruthenium hydroxide, ; 


Ru(OH)s, soluble in acids but insoluble in alkalies. 
4. KCNS, in the absence of other platinum metals, produces 


gradually a red, then a purple, and on warming a violet, coloration (very _ 


characteristic). 

5. KNO: imparts an orange-yellow color to the solution, owing 
to the formation of Ks{Ru(NO2).6] becoming a beautiful dark red on 
the addition of a little colorless. ammonium sulfide; on adding more 
ammonium sulfide, brown ruthenium sulfide is precipitated. 


6. Zinc at first colors the solution of the chloride azure blue, but — 


subsequently the solution is decolorized and ruthenium itself is pre- 
cipitated. 


7. Hydroxylamine reduces ruthenium tetrachloride to ruthenium | 


trichloride (difference from platinum), 


IRIDIUM, Ir. At. Wt. 193.1. Sp. Gr. 22.4. M. Pt.=2300°? 


Properties —When produced by the ignition of iridium ammonium 
chloride, it is obtained in the form of a gray, spongy mass, very dif- 
ficultly soluble in aqua regia. After being strongly ignited, it is almost 
completely insoluble in aqua regia. | 


It is more soluble in aqua regia after it has been precipitated from | 


solutions in a very finely divided form by means of formic acid, or 
when it is alloyed with other metals (Au, Ag). The metal is unaffected 
by fusing with potassium pyrosulfate (difference from rhodium). It is 
oxodized by fusing with NaOH and KNOs in a silver crucible, but the 
compound formed (Ir2O3 combined with sodium) is only partly soluble 
in water. If the melt is treated with aqua regia, however, a dark-red 
solution of Nag[IrCl's will be obtained. 

By heating the metal with NaCl in a current of chlorine, Nag[IrCle] 
is readily obtained. 

Iridium forms the following oxides: 


Ir203,? IrOz, and the hydroxide Ir(OH)a. 
Bluish black Needles with a Indigo-blue 
metallic luster powder 


The dark color of the chlorides is very characteristic: 
IrCl3,* IrCh, 


Dark green and brown : Black 





* According to W. Patmagr, Z. anorg. Chem., 10, 322-326 (1895), IrCls and 


its double salts exist in two modifications: dark green nied brown. 





. - ‘ & 
ee Eee ee ae 





THE PLATINUM METALS 613 


REACTIONS IN THE WET WAY 
Use a solution of Na[IrC],]. 


1. HeS at first decolorizes the solution, owing to the reduction 
of the tetrachloride to the trichloride, accompanied by the deposition 
of sulfur; subsequently brown Ir2S3 is precipitated, readily soluble 
in (NH4)28. 

2. (NH4)2S precipitates the same compound. 

3. NaOH, on being added to the solution, changes the color from 
dark red to green; on warming the solution it is at first colored reddish 
and finally azure blue: 


2IrCl4a+2Na0H = 2IrCI3+ NaCl+H20+Na0Cl. 


If the solution is now acidified with HCl, a little alcohol * added 
and then some KCl, there will be no precipitation, because the Ks[IrClg| 
formed is readily soluble in water and in KCI solution (difference from 
platinum). 

4. KCl precipitates brownish-black potassium iridium chloride, 
Ko[IrCle], insoluble in KCl and in alcohol, difficultly soluble in water. 

5. NH,Cl precipitates dark-red ammonium iridium chloride, 
(NH4)2[IrCle], insoluble in a saturated solution of NH4Cl. 

6. Reducing Agents usually change the solution to a greenish 
color, owing to the reduction of the tetrachloride to trichloride; or the 
solution is decolorized and the black, finely divided metal is deposited. 


Thus if the solution is warmed with KNO,, an olive-green coloration is 


produced: 
IrCl,+KNO, =IrCl;+KCI+NO, 1. 


If the solution is boiled for some time with an excess of KNOs, it becomes 
yellow, and a part of the iridium separates out in the form of a yellowish-white 
precipitate, difficultly soluble in cold hydrochloric acid or in boiling water. 
The precipitate has the following composition: 3K;[Ir(NOz).|-K{IrCl.]. 

Oxalic acid, ferrous sulfate, stannous chloride, and hydroxylamine reduce 
the tetrachloride to trichloride. Zinc reduces it to metal, and so does formic 
acid on warming in the presence of ammonium acetate. If considerable 
mineral acid is present the reduction takes place less readily. 

7. Chlorine. If chlorine is conducted into a dilute solution of 
iridium tetrachloride, the latter becomes reddish-violet according to 
W. Palmaer,{ after some time the red color disappears and the solution 


turns brown. The same red color has been observed by Foerster ¢ 





* The alcohol reduces NaOCl to NaCl. 
+ Z. anorg. Chem., 10 (1895), 358. 
t Z. Electrochem., 10, 715, 





REACTIONS OF SOME OF THE RARER METALS: 


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THE PLATINUM METALS 


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516 REACTIONS OF SOME OF THE RARER METALS 


at an iridium anode in sulfuric acid. According to Palmaer, the red 
color is due to the formation of an iridium compound having a higher _ 
valence than four. 


REACTIONS IN THE DRY WAY 


On being fused with soda, the gray, brittle metal can be obtained 


by the action of the upper reducing flame. It is insoluble in aqua — 3 
regia. 


Separation of the Platinum Metals 


The separation of the platinum metals is one of the most difficult 


tasks met with in analytical chemistry. If the metals are already in 
solution, the table on p. 514 can be used to advantage. If, however, 
the metals are present in a more compact form, it is quite difficult to 
bring them into solution. 


In the latter case treat the metal in as finely divided condition as possible 
(filings, etc.) with aqua regia; this serves to dissolve the greater part of the 


platinum and palladium, as well as small amounts of rhodium and iridium. 


Dry the residue (osmium, ruthenium, rhodium, iridium, and small amounts 
of platinum and palladium), place it in a porcelain crucible and fuse for some 
time with ten times as much zinc (or lead) in a current of illuminating-gas.* 
In this way the platinum metals are alloyed with the zinc. Allow the mass 
to cool in the stream of illuminating gas, and treat it with hydrochloric acid 
to dissolve out the zinc; this leaves the platinum metals behind in a finely 
divided state. Filter them off from the acid solution, dry and introduce them 
into a porcelain boat; place the boat in a tube made of difficultly fusible glass, 
and heat to dark redness in a stream of oxygen. 

The greater part of the osmium escapes as osmium tetroxide; it is absorbed 
by the caustic soda solution and tested according to the table. Mix the 
residue intimately with sodium chloride and heat in a stream of moist chlorine 
gas. Dissolve the mass in water and examine it according to the table. 

_ If lead were used instead of zinc in the above procedure, the alloy should 
be treated with dilute nitric acid, which dissolves the Icad and the greater 
part of the palladium. Precipitate the lead with the calculated amount of 
sulfuric acid, and test the filtrate for palladium by transforming it into palla- 
dosamine chloride, and then into palladium cyanide. Treat the residue from 
the nitric acid treatment in the same way as when zinc is used.f 





* The operation can be very conveniently performed in a common clay pipe. 
The gas is conducted through the stem and ignited at the bowl. In this way the 
gas continually streams through the molten alloy, keeps it well stirred, and thereby 
yields a uniform alloy. ! 

+ For more detailed directions for separating the platinum metals, consult 
the work of SarnrTE-CLArRE-DEVILLE, Drsray, and Stas: “ Procés verbaux du 
comité internat. des poids et mesures,’”’ 1877-1878 and 1879. 





A 
PAGE 
NM ON ce tide ve 4a Sloe aieaig << sie elspa wt Rove evento? 122 
Acetate ion, effect on ionization of acetic acid.................. 0. eee eee 46 
IE MS cg 2t Ree Ih Sens alc Fla mow a's pig. «a gins Sut Sea's 00g ao are's a 341 
I ek er 2s e's aia gia 6 Sa sist k costae oe ae ti elSic ove eg ka ead eos aes 341 
jonization in presence of acetates................0-2 eet e eee eces 46 
RIE a5 8 ory 'ar'a ore aarula aks WE o sk ad Rita ad male a talons < 
IE eA EA ST oe a in EK Wan aah 2% aye ee sun wid ae aoe 80, 342 
Acid properties, cause of.................. Soy So ale O70 AO Wea ears Lc pee ede 5 SST 2 
RU OT ew eA) Diani gare hee big ga" aie 'b4'-dlemp 54, 447, 455 
ewe, NINE NNSEMIGD chao cis 0. a gies ae pls al Wale © al digs ss 24. doug si eR prs e abe Sie 7 
SSO RIS SPSS: pan aS a CAD rae te Pot NOR OO ae AE A a 57 
SO SARA SASS ee eA RT aR Ra ig 284 
OEE, ECR PES AE IS Ihre JT SR Ur Ee SUT EA OE Pe 56 
examination for....... Sh UU. TS ROW ad tara aa eee te gE 284, 446 
hydrogen of........ BA te is tat ergs ey emetic Wesley ROSE Seo. hee 32 Wo aA eS 2 
REM Fer Eis oe deer oa ee Se ee Sa 8, 10 
NE GUS Se SS oS BAT | Se PR king age AaB CORRES PoP tne Sy eR PTE 6 
preparation of solution for analysis of... ............ es cece cece e eee 447 
URGING CONCORD UTOLIONS 2% 5 oi6 os oars 8 oS hips Seaweed a5 Pare Saye NEADS ak 71, 72 
MMT aRER VCORUCCEION, OF 65-6 6.5 oslo x sos 0 Llc deld whe a die cae theleeg bb as mie 54 
EE eat ER RN ie ar Re See Gea eS PERE 3 
ES ea eer reer ORE es Seen eae eter eae 28 
I SEESIIS RAD TR Sr ECS Se an Pa a RP or Poa 19 
AT SG yc a aici iu, wici'd kak Sete w cal Ca urs Se o's v oR Re veld @aralwale ees 413 
Tl a a ieee ere SG eaRN SALARIES wh Saaz Ditsy 2° haie toed Were ae 82 
MeNIRETITLE OF ALU EL yo oss Seis kei Helse fo ainke eursa ccd yes 0k delela leeis 105 
Re soy co cidib a viv. 8 dal'ary, 6 5 Wea sce vad & io ew Oaiey Wowblgws EEN 77, 456 
MERI atric a eR NC ees sales: Mas 86 ee aoa Eee oias wees 97, 459 
III RORE TONNE O02 6 oa «f'n. v OG 4 vce bi sv RUSS Es tosvacies Wadee Reb 72 
MINE BLLICHL URS J 524 Sin Was sds we 8S dy bare ae S Cao's oN gk een 419, 460 
EN G3 Vie Sik Se PAW GON Ca dhe ee OT Ou sls ode a ck hte MOTE) Kee 77, 456 
EEN SPO INMENOSIUM 5 6X occ cs ek loess we he Ho Renae ewes enaels 96 
ROOD REN cn Bad SA ICA gik V2 xb didie se hig eed + Rea Rea eee 188 
Rei a gin gig 0. ct cin kc See e ew de ae ach Wied Beale dade epbes 101 
IN EI orn S wie 5k speck e WAGES clei vied dle in Xu AS aa ONS 111 
Goparation tromy alkalies. i 2.68. i co iels cece suey 111, 113 
MUON ok pissa oa CR aM Ais ee OS Ke 188 
RTM SS ne Clara avid pa ask bos 6.6 ora HAS A eee kaw arpa we 349 
REE CEOS Ws ose oi oe Fee Nn as cen eeee 391 
PONCE NEA Oe ORG cas SOT TCRSY 90> dew oe Saber on wi 48, 447, 452 


518 INDEX 





Alkaline solutions, notes concerning.......... Raber x ie Siege Sm .o.- 488, 452 | 
PTIATIU G6 4'og wih oon 8 ok PRO aw Vila tice heal, Cea Rea aie Bia era rena a .. 468 
Allotropic ra0dificd tions. «0:85.54 swan cae dapidin. ald Mole Sale Wh a eee Selene 196, 339. 
Allo Ve ARAL VGIS: OL. 6: oe 5. < oield.n sip hus Biinas Kiateias pl-nuaerac tee aaa ke he 
Alm, 5. 5 55s won Soe eb 2 plvonlb yie/h 4 le allege eR wee Wamls, 5 eee WSS ak 
detection in the presence of organic substances.................. 
oxide, method of dissolving «0.0/0. 5.6.0.0. Shene > 6s alee oe 
hydroxide, ionization Of. .... 05.25. ..s4.« tia eS b> 3's oe eee ee z 
separation from other members of Group III............... 189, 443 
PEAR eo, CSS a vig poate Sle crass wis gies & salen Sy oe aS 25 
PATGUNGBE, sisi soles as aves bong 9 6 6ieiwie,c BoM Aw # ol ebe oh © oats) ele-p aimne Stain a 
Ammonia. See Ammonium. 
equilibrium with water. 6.6.00. sso ves eee sos ds akce et en . 
Mi rinking WAGE. oO Sis sie ews Masvimns Oe Adds, oo we 2 ee 
Ammonium carbamate................05.eee eee PS aS peas as 6 We 3 
GRED UDB. Sock waste ek ae we oh LN OR posed ee a Pe 72, 103 — 
chloride, use of......... EPUR ERDAS NAAR RES Do 19, 94 — 
hydroxide, ionization of. oo. 5 5.5 ou. Woke etn ees pe 10, 18, 46 | 
ions, effect on ionization of ammonium hydroxide......... 5 ane 46 
molybdate -reawent 605 Fae ie Ve tl ee 72, 231, 379 
reactions of... .. GAs acs als tira ara bye as wattage ale kit aha eS . iheegene st 
salts, action of strong bases on. ......). 6. 6) eee ee . 87 q 
effect on precipitation of magnesium.................. vo. 94S 
BERTUT GRO OF 6 oe, sy ats ole Bee Slew es ate eal 5 we eel ao ee . 92,98 — 
ealide Brouiie si EA #4 as piaide.b\h e o4.4 a babe aoe een oe 461 
> BVI OE. 8555 enh canes pe se 189, 443 _ 
Amphibole group of minerals...) 2...) eo nid citamiele eon oe lee 8 ago 
Amphoteric electrolytes........0....ce eee cseewecees ate ea eu oh et nn , 
POR NNT GS 5 5 a. 5 cin soi hwo Wig coded nln’ p's ere thie 675 050g WAL Lemus teRIS oc 
ial yates; ‘Chemical: 3.5). :3.5...5 sks cccaly sw ere be ce ve iv ols aD aleve aly Sate cheer 
efuialata tive Sc 506s NAR ST yolk cw ers, 5b eae ee 
CPUAMTIEALEVE bosses Sie este ATE. SEN Win aln k ie Slot a See 
SPOCUTORCOTIIC 5 255 eos Bis oe ee he's's 0,0 pla bie lgps) «waa late’ Bee pipet en 
HB DOMIB TAG. iw oe ios SS ewe 858 BS ale wee wis me ore A VND pha nn Wie ee 
tt eee Pa ee Rac aeRO Dee 
Anglesite........ AGES: SCUBA OE ESP aR 102, 205 
Pat a nS Ct 
SIND sis ais ES oes we ele ae oe ed ety Ws Se Ce ORs Se Port. 
examination. for. 220i. cb5 RRS OPE I Pe oa 
prelimimary examination for: .... 0.60 525.0005 LALO oe 
renotions Of 60/50) ev eee SIS ANG Fae esto ae 4 
Mmaboraites 25k ieee ab eas ORES eM ap ae 
UIE. 25 02S SREB OE UNA She OA ov lee WS ola 102 
ABUMODIc COMpOUNGS s 1.15.66 i i SWB ES WO i et 
Antamonous compounds 26065065 0 i Le OCA a 24 
PR SATAIOIAY 6 aos 0 5a 88 OREN 8 TAs OT RIE OMIA eben OA respi 
separation from other metals of H»S Group............. \, 272, ann Se 
ot) | re ee Pew wL rewire ewe ere ey 
NEARED? 5555-50 Dood Boba PERE pe: 5. od oa ee OS Aa ey ee eae 102, 205, 2h 
UR TORIB 6 oicecic candies es Claeys coedaunceeiPieted bv cree e epee #9 


INDEX 519 
PAGE 
MEO MRD ee epee OVI ee eC L Coy eee ETN ETE N OE TE RU Ca we beige bre 102 
Ss 2 GER OS bs A RE aS a ak ine, PR gD La nena 278 
UNDP SUEY REIT A. CRS Si vs lores evo d Rye ew kage eeeud ase erp wee 4, 5, 7 
RE REDD eer LEAL 6 yo2 ip FOR RPO ers: c-u/ata eh F oha dodo vieiira sd bie ge epbinleee water’ 223 
ONE AE De Oe ep hard at nn oR Ee Dey RE ae era 229 
RMR LORE TOE (FP hc OC ia here a ene 5 cated) va ee noe Obevis Ohne 228 
detection in urine, blood, milk, beer, etc. .............00 cece eee ees 238 
GREE ADOT vic i ciesccadesy buna oes Re Dae Ne Buy eae & 236, 240 
NS Pe DD Ee or ar ge ea a fe SAE PE >. 1288 
RIES PVs cr Re DO IORI a OOo. Ls Ape ee se Pe ewe 232 
EN ESS evn ie eamren ly ek aia e Cats g0 ba Dae eb a eRe 229 
Reinsch test for......... (1 ats, OEM mild Tatice RAUL eR ee Merges Oe ae ore ESD 240 
NPERIPITRLL re Vic tic tat roil oo ol arco Ci ariph hint oh ha w/e Uo wc ave See av ed ON IPT RINE el 224 
IIINPIREOHAT), GNAUV SIS OF... 50/6553 woe oo arn shelled Selle tie een op Sorel biie gk GT es 275 
NTE BER IIGLONLI gs 05'p2 +51 os nese sist ap ekg eps deal al UU Tp ieee lu hel dod ete SOON! ghee 224 
EET St be ls biter ok ee al « Cc Wn A Why oA wo RAB Re Oe Wee 223 
EN bse S Si rie TGs tes oa ate MAMIE aie Sipe whe Wm MT wie sw RASS ered wale eet 224 
Arse (arseniuretted hydrogen)... 02.52... cece eee eee Oat veces 232 
TE Nero ey ea eae en 7h Seat, jelabe neh oe oe fe cn poeta severe sere GS Mo 93 
NE aN 56 Fo i er vg wg Saga vc tote HED, thw dy wien OME Re Mee uttbe ne se cere le cr taee te 214 
IIE nated Plies Sac dale ST sg Gk ena te lyin tm is Ho, candela 4 ve whales “Mle Ss 93 

Autenrieth and Windaus, separation of sulfurous, thiosulfuric and hydrosul- 
. IEE yo Ss eee i Savy hae tie Se eae Se Dis Wael. alae cig win Oe See 390 
UMMM 0 28 che aac oes uss ss dnese oar tna pNNS ew ass Se ee ey Ke ya ee ey . 170 
TT a. ah es 21 sae -g oy a dna. sahara ond cota ba te tn vou? lle) 0% ou. eR Sape ed Aes Wha s REN 214 

B 
EM SSO Coal ie oe atid og tweeters a Pele ete AGS eat wy a Fd SUL dee ee 462 
IR Na hed. Scie oe eet ns SRM iol eh ETN inl as ee + Ge EN bo Vee 102, 108 
EUR ST ees, AAS Fy 8a oe hal) fh GU 2 OHH a goa bon > RIERA A a UAE eee 108 
Ea ha a ETE ie Sip Rta AALS By SQ CRE Eee Sy Rie ER Map ays ese 2 8 450 
separation from calcium and strontium. ...................000% 110, 113 
ETE RLOMSOOTPTIM MMU LOTS ONE oo), ise! 6 ua san loch Su. @20! olere) pala phe 109, 437 
Bases. See alkalies and cations. 

ERS ARREARS (ae, RAE on ete a Re lig ean 8 REL ia e 72 
RENE SOR SRA TE MMe 2 ARP SESS Re Zee we IAL RONMNE BD Se o-E De ego 54 
ER NONE st C1 CRA Ne ANT Pea Fe APM RR Se SMe es 2 
a at holy Gey ped clte es Web Ra ahi go I etal 10 
IR NE 2s) Saas gocts Ny aa the ah DEM Cera een do eueaw RE Sane ee 6 
RR YE a Ly bpco edie f lfteterh ceed tor Sawlett 2 
UND SS oy a pie led p bbb wiv bid cle op HUlpL FNS DUN wrk 125 
ESAS RICE er CaN are a re 64, 425 
near seem RT: GEIOTIIG: TUS, Sg ng vc vp oy oot sio'y t's pee ab pap a veeade eee 238 
Ie a Te a et che aah ete eae avinte Ge wens 310 
Ne NTE nN ee es aaleer ak Pa-aue 6 ba blarted meee 461 
Ny PR Re he Nh Fe Oe aan wlohe a ide veh Coe rey ee re 8 461 
ICOM TERIINE ER ol go) hE eg ule iy pv ua s 485, 496 
Oe ee ey ee ON eer OE i ro 232 
PeGhUcritetn ee MIEN TORRES re ok he. ay ecg y gene bee otae 228 


520 INDEX 
Pisnuth, peactiona. +6266 asc choc seh owe on entse Ob eure a we 1 Se 
separation from other metals of H.S Group................... 272, 4 41 
pabastrale. vias ak isa ae STE es +6 chen Said gales eee pe ad 
Bismuthinites 2 .:..50.5 Fs eva bh daw olin we boule Shean ee a Wee ge ee ee 
Bismitithalc ACI)... sso abs deeds ese e aes ya 
Bismuth-sodium thiosulfate reaction................ cece sees eecescecees 
BORDA as 5 5s os ital we as so ioe als eee iy Da ae were 
Bivalent acids, ionization Of; oo... 6 ses a Vek a es ca sem os eae 
Blochmann, strength of reagents... <-.. 0.6. 6 NG aol oc = 
Blondlot-Dusart test for phosphorus... 2... 20:3... .s0s ees se e003 ee eae oie 
Blood, detection of arseiiie in... 0.60 bee de ey ane ds dene eee ee -. 
Blowpipe reactions. 6. i s.0 a ee oh en a Paw eee es Oe 68, 426 . 
PR eo 5a 8s a ae lo ate ee eH pew thle ON io ink Pee a sepa o> 25 
PROEACID DOIG. Co 'h6 oS 5e s ae ee hw ko She Ree bee ee eae ev ade ce ence 
Borates, solubility Of 5.054. 62 Sia 6 Saas Sats sen bo ptace ce iss a ek we 
PROMI Pe G8 oh goo Taco k 0 CA le a eh OEE, Meas ee 82, 357 
BMAD on, oss cages goatee Me wesw kerk baie Bes ere eee ae ea 64, 425 
SE BOUL. oy ac 5 o koe eee ae Uk) A PhO Oe aE ee ee ae oo sree yin 
WPRRIGG 2 Ne a ns Bg oS Daa ee ele gw die + © apd aaraly ileal dee fia ae ann 
Breakage.of dishes, cause of, . 0s) vs ov. oes o's de oe 08 4.0 bn oa gk 
PRTORIUITE. 5 50 FSA Re Seok aa os 8 we ah nes pb seary 9b Dt ae 
Bromine: (free); reactions, . i666). Ss a os oo tote Saws + a > 9a 
detection in non-electrolytes. 0.00.06. 0c ccc ceseencesocceses Sen 
presence of Cl and Tow... sc ces es late By ceed Se 
BOTOORTOBE So Sa 2 05,3 SENS coca 26-8 Male 8 lp rin ada Rela, LAPP Os eta 
BRCUUE TORUGIONL: 5 5.5. isis) o<cheoeccBe ppv! atnle are cop FE AE koe oer 333, 395 
oo ost) i a Onc eee eel MME DEINE ER ara gtr ten 
SOIOWIES ids ob bo 5s agi TE aie bw Swe Obs 4 eae ae Cee PR ae og ale 
WATE, ches ois SSRN SS Ma a uk Tetwike bw Woe owe Re a Oe ete 116, 284 
Mame, parts-Ols J. 5 sie eo on FG mw bb ob eee else dee 8 Da 
POMBE. SS ei ie eos oO aa ee Pie a a ee wk we vie Gee eC Cen el ee 
C 
Cacodyl oxide......... END kG ine bs we a Ge SD bs Sen eee <i 
CML TORCHONE. ook cools 56 b9:a ee ead ee oss bb es oe 
separation from other metals of H2S group.................. 272, 441 
ammonium sulfide group............cccesecveecs : 
Omamim, Teactions?.. 0.5. ova c' os bx srs.op0 > vive bio biped alg a 2 be dl ee 
separation from rubidium and lithium...................00e0eeeee 
MSNA 66a a5 5 a0 Sigla pine le 6 4 U oo pro Noela,ale idle, ¥ Rye Sieh os Sakae te eed 
BT RR Se Pe oh Oe Se Stee 
IN ob gown 5. bhi e ten LAM acato- te shih 0: bcos oilaghe pew LS 1k Sa ahh lg eRe ao 
RBI, BOACHIONS. | ..3 5 bew.ow So oy e's 0 seo Veoy aya rae bees eee a ee Tunas 
DATISNAR Ap oC cicecabed thin bin! s, sracach coe bes Loa habeeca eh kee atte eae ge 105, 
WGP oo oe FS GR iw inn oin eo Jie one wip nas Oo tele ee ae 
phosphate, solution i0:ACId8.%,.. .0 + is ecce ceed e vy oe vic nls 0) SHR 
PROSDEMIE. 23s idk adn aaa biganseehe¥s eek se el bie 106, 439 
separation from Ba and Sri... secs .c wis ss hb 0c Ue suiels oe sae 
ROMEOION ois 4:4 aishe v0) Holmen besten. ovo wlaia bese Oo tint © DR eae ees Me oe 





INDEX 521 
PAGE 
MRI MIWM: ORUGION. OFF Uvshs sc ccs oe tc ccc cee ctcs seeders eens Sit sa oae 439 
Rane MTEID. COLOCUION OF o..04.5.)0 cole se eicieleeca's cc tccueedreewieecicessed 220 
SrERONLeE PCHAVION ON HMNIUION.. .°.°.%'. c/s a's bin elec dive peelaneceeecetserecce 355 
Palapiity Of oS ie ee a BP? Dare N IN Ere ie te 353 
DerMnnNE Men SORITAGIOT OL... 2. ss wa 2s s eosin ep ooo ss cbeeeetteneses 8, 10 
Oe = Ie ae as Sea aon mip ahr) Meret aD ONE Yq REET Si 352 
Sere Sc ee Cea gS ein SV CE Cw tol thee ew eiles oe 421, 436 
SE ga rar or gee ea Re eae? Ree 78, 93 
Gimpociation Of... week akties PML EO eeu ce os oats 16 
Semen OeeCIION. OF POLASSIUMD. . .. 6. eel ak ee ee ete eels 80 
III Pa or Poetic ie oh s Peas eR tho d's Le hades Bless eta be 487 
REE. oO. eee Ss Pek PS Soot alec ghee anes 403 
ET ges eas al Suc wSlag ns Dlg oe hehe tg oi alalg so 249, 437, 462 
Cassius, purple of........... PARRY st CLAS ETS ROO N at ces es oes 263, 265 
MOMMIES ee oes a ce occa es cdot ae eRe Re Tea ee city eWLee eee ae 5, 10 
ERs rs a SOR I Tea Site td Ce ogee sees ea eeudeas 5, 10 
SINE Soy sos au Die 0 cals soe noe wis Wea ¥idlee Si Na Win tg Vinars tals aie 77, 440 
Caustic alkali, detection in presence of carbonates..................002 e000 ' 455 
EE Rt PO ee ao 18 Se UNO R eee de SU a oe sels bee es 102, 106 
RemrIeONTIOUTCS “FORCUIONS 62... 21k 5s ek ek a cate ete ee eb e eet eenene cae 470 
MEMES ry. cla ele Gas ocd, satee oGjd kde 515 wend Sere eb Ole wy a ee 478 
SRR Se S Te Noel ees UN a GA's a ek ARS oe ate 5 be adie Uae 469 
re PE SEN ea FY Sia eel coc wielelaierete’o esi utelaws «t'ee oe cea 469 
Suman TOONTICtINCls, PEROUIONS) 22 62.2 kd oe Sb oe dice Loa ge he ede e se ots eee we 469 
RE SIE Los Siw! ol el hw Sal vhs Sto Kw Wf dela o a nvoralele eyelet ale 102, 205 
rr AS a Pe x aie Sa oe Su scone wah soc pix ahd. toh oc SEE onln eRe w vised 413 
EERE ag SEU Scone arn a Cae ee RE Py ones? BEEN S| 
RE eT ek hs Dar inthe Dal alata wi alate alg’ gtaitiateriigj Pod Pawan 0 ER 205 
Charcoal reductions before the blowpipe..................502 eee eee eee 68, 426 
IE PERRIN ot fF 7,4, Esk o vke eB ose uavs Uwe ly wl Nee. alee A ated «6 65, 435 
tar iar 28 ee AEE ORS eae med we ee ead cae cehede es weae s 214 
Chalk. .... RT 1 AO eyed hse) Lube Taw acs Ace Mette eaten 102 
TS EIR 7 en ora 2 SOP ay I PRM a ert eae 1 
IMEEM TOERNA S os) ou tne tac cig’ v6 a Sick Ton’ vid Oe caida ata nod d'é.g'ww ge arnree 13 
eres et ed Gillen Dy Sk ev ag bt a Dees eld eae aie 3 Ord wiataieie «ere 1 
5 LU. idk Ay ck asd atthe a ae ands ofp lu's v0 wb ev oe Pate e aelw nee 82 
I ras Scie gin Neco gis a ahh agave este Se aaie o cidlaleaig’e slaidieraty dace Mad 172 
8A rg rahg. sgl an grePanat baie oes 4 GE NIV Wide as Sate Bele Ba 397 
ERMINE oh oy hho ois. ab oka woes eials’ AelaaL ew. cha’ nl ed ¥)0 ¥ PeeWee eo A Ble werd 397 
detection in the presence of HNO; and HCl................... 398 

Chioride @etection in thiocyanates. 2.66... cek eee eee ccs nee ece eee ea tue 326 . 
RJOMOTIIGE EHAVIO’ OF UNION ic. bs. etd ceeds cee ee cc eet saws eta paee shel 290 
EM SRL K LG) 5200 5% 05.) Sed of heed Seach y! st ah shainte a shin Gy voala: the Whe Wig a 5 287 
ea INET MEMOS Sco Sg 4a aie 4sar so ie when wis a sleeve voy Beedle ww hes 291 
SALEEUIOM IN NON-ClECEPOlY LEB. ios Sod cree ev oie ee a ee ce meee eee leas 290 
TPOMOUIOR ETRE ENCE 6 ose oa ho se 6 sles OR eager a wis ee we om 306 
NARI Ane arity sletdiale Slo & S's eS lek us See Spe «OE 416 
CHloropl Atte AGH (GORENG ele o eA ies hae bee ve see rete seed epee 73, 269 
AAP ORUUM RUIN Ae talcta ps nee vies Stores os Se PW IM ate ne 6 0 Vole Fv We ONS 8h ee OEE MORN OS 135 





§22 INDEX | : 
Chromic acid, oxidation by means of... . 21... .. eee eee eee ce tere teees 32, 138 
eompounds; reactions: . 2.0)... 6 Rk Se ese eae +o 
Diroauts ee eer eae PEP re te sae ANRC . 182 
rig lyeri tk SF ie vicky Sex keel Odeo acs bmiteia wahleaeee Ohad pe 4 
NURI in 5 Pas Sa kd odin wi Ca viewed ep RR TENG cain ga 32 
X20) <0: (ee 8 I ee ote aie eae ag aie ae she eee 182, 4 a 
separation of other members of Group III.............. 189, 192, 
Chromophor er Perr it ORE Ser IAN re ES 
Chromous compounds, reactionsS...........:00ceceeecuceccceussceeunee 132, 4 436 
SPAT yeRIOR GL oy as ela Sinieig cede c's od cco deo 6 ob pale Uleuible o cneeee tie ean : 
de En ogee eR POMPE AM Gs bei yE eC Le ori. (195 
Ortrates, solubility of: 55.0618. ee SCE asec kbc duels va CON ee 
ATION occ Six, Sas oa i ie Cie pels, son 3. ca awe os oe oe | 
BO | ee Sms SM Cee Mea OSE TAU A TGRE MEMES SM 
RABMGCANIGG ooo. scinropacy, eoleie. boas hc ¢ ohcalond wodie gro. dsm Wongacn ile maiko Oba uae 
A es a eh Nav Wb res cresan ad hemes ARK Ua esl Co ak Day ee oe 
Closed-tube test......... Hei ee eOR oR hed Bk Die Nee URN Rane oan ea 64, 67, 423 
Cobalt,:complex ions of... 2.4... i sees ie eee eee Nya, Se ole eee see eae : 
Goeobtion iy nickel alts. (525 305 cn cienic = ohn + aise hes Aoloas - 0.3 9! 0/, 184 at 
hydroxide, acid properties of .............20cceeeeeeeeeees oes 180 
i. itt) A eee ee SOMME a Um SIGKIIS 178 . 
separation from other members of Group III............... 189, 192, 443 — 
SPMRIOARRCTUMANI GPU... 5 sc ou bat owen Roos Moe hao Pee eR NTT sate 181, 328 
AON GILG G) 0.5 aie soo; <ieco.es Wp ee Ne PN PR TA ON EA) eR Gs) be se ajalais oases a 17S 
Codein test for selenitm .... 2.3.6. odsi eed ecsscs sues seclawoe cena anne 499 7 : 
MONA aL IAL IOTAE 85. 525, C niyo 013 sic yes oid a bele's piv elec ere geolelend eto ea 58 , 
POOCMDCRH 0 5 gk CAs Roaleia a a ee aes 58, 127, 218, 415, 495 
MEMMRMAA Co rsits: $< ooo: Soh np lw wn eb ol opp mane bm; Sup hanl Hes REA alk tatoos pe 59 = 
oe i a eT ATOR eat mre Se 60 
Color imparted to the flame......... 0.6... ++. eee e eee eee eee ee eh y tas 64, 427 — 
MAMMIUAUAONO 5555. aio gb ve acc v6.0 's jaaete op tis wip» Petia. 9 oss eld Siena So ela oo ee 499 | 
TNMMNEM RCC we vulid sv ie: co poo, vinta Voir Wiwtohe soem Sb (kot Me vo RST ete 480 — 
OMT FON OFS Cb os sok DOSS os ied ws wera tw Bhp eo ware atl OO 45 
COMATIOR TOTS. 25 ode kcoicicie 8 dip & ods eerie ¢ 0 di8'e.p wae Vale epee ea ts emus Sn 24 
MEMNSOE OR: SN rsia Jay fin eee re oan 254 
ealta'and double salts. 020: 6046.5 60 soe Bed soe 20s Oey oe 25 | 
RIMMMCRLCEAUION 56655 fv oko ovis bolas vg so dies nlp ocd Had ov a a phip wale ile 9. 
changes and ionization............. aca-0 0% blk. pele eo ae 18 _— 
mole i Sy Soa deren vie bin 8-9 8 0 wie Shia eb Wel ghar 9 
TMOTTENA  o5 5 so na'nce+ pele 4 © einipee eo 9/ore lp binlleee we elegant 9 = 
Of TORMHUEE, 5. i ies) c.0's + a nwe.n-o Gi8Se win} we ad ee Pe 
Conductivity and ion concentration. .. 252. 660. essen becky egestas enya 83 
Conductors of electricity... cis. .i6000'. sy 6 4cs pele oes 00's s 408 trie 3° 
Constant boiling acids, 5 5 6\)be5sc's4-0 + + Jovot ove ¥e be ¥ 0 Unicel 57 
NS S57.) bs co aie ac Fk Vd aR eh pws Soe Le eae io 46 2 ae 
separation from other metals of H»S group...............+. 272, 273.441 
(NEL OHS Group 5... 6-03fa'os a. eater Fe ck ae weve BIZ 
A RMINNBL Sooo ig v5 Konus! -9 4's igholoverWiaCdre's 6 4,ug ala kesh op ean dialy oh en 125, 438 
MPN RU ORCA 8 as. od oy on oe hie Up ive Sf eve ple toe ea a ews gtbr ie ol ona 362 
MNO EOFs 0d slo iad Wea the CEs Dame Sev uae Phas Pia ee ree e ee .. 132, 205 — 


PAGE 

Crookesite .....6..+.006- ES ES OPER SCOR Cha Id Sareea th a a gee RPO LID ee 485 
AMAR irre Tan gd Sethe Sin cals «vos e-u's atk «5'9 vaio KKs 89's PA ie ge rare PEEP 82, 125 
NE ro aly seid ies sca se oy bn ands Ves BPNRe eee ee cers 152, 160 
NIrILOUNTTONITICLE, TODOUODS 6 62s. 6) sacs esc es teed sa eh c6 aie su bend die wore 217 
ERs e te Cn ta aN so ootny Ag ACRE Os DEAE Sek Re Siem bleee Heke 220 
eR bie age bik. Nak bie cao le Veu usc ees ka cede Rae Aes ae oe 214 
NM gata ici oy asp. shee yy 4 dies 'e be fv d.92a'9's bine Lime CR ATER N fe guna 494 
PMI OOMIUIOUNGS, TOACIONS. ,.. 0/0 cose cc dace tess tedvecceevccedsvdues ear 215 
DRE ste wna a psn Sie sees « Oo Pei ohh eid Gide lbmrdiel gi xa. e ce cree 220 
ENERGON Sy ee ek wd Slyia wae Eh hs 0 8 tp ow ieus Fads hv aes Sree woe 342 
I eee en Sino Wich ne ae «wile Oe 6 ghia ae Si Kia'seg Wale wale ee oid wD, He 342 
EES a Baer, ee Ot eee Ae arn oe Wt ree Simmer a gas Gry 311 
PERM NES SOME G UIE olla lc 5 1- ciag See aca Wiebe ode Dio a bd CO odin Oe yep 315 

complex compounds, 24, 146, 181, 188, 219, 280, 312, 315, 318, 321, 328, 438 
decomposition of. ............. ME ee L DA: kee CEA OE Uae 152, 322, 438 
Cyanogen, complex compounds of, 24, 146, 181, 188, 219, 280, 312, 315, 318, 321, 
Z 328, 438 

compounds, decomposition of................4. 152, 315, 320, 322, 438 

D 

eT vo ca oo ae Cay HR Gata Ck AEA KU db ATE 6 Raa 42 
a EMU EIUIUNTACUS TORE dhe cy 5 sh cicis vile bn on Gia aD ev eb o0ik him o Seed 418 
MN MMIII og Os ake wy oy la oihckS hg 8 Ree Potwine mv es hoe ee 6 Pie ie 367 
EMM A Sek. way eae oF whee. Spb rate WO DUNS ele bd hin se aeue 34 
IE RTI tastier aoa. 5 la vie 8 ak dana ai wie CRG B-o-4.6 © lnceyn ish aee Bla gee deen 516 
TIEN He ciel vata lc o's ahi cies woh ev aft © oe Maes ROK AA RTT ee wae yes 125, 142 
PIAL. ACT ATE DOGIOD Of. oy ics neve chakecdbecev des atuwdetebee 32, 138 
IMM ial cary iso he Gok We KS wip View a pono) boa, Sw Raiecln 6 SS te Se eat 316 
IE Reg gra sslu fas ald. V6.3 alow 04/0. 04 & bio apa. wie Wey pine oars Con Ou aie bk 469, 473 
Diphenylamine, reagent. ...........e0..c2eeee- Pte Eeyore eee 332, 394 
RIMINI Ce 5d sre ae go Hic anole VL oe we ek SMe NWR ee CE Mee hoa oe 177 
SIN PRO OR ULE: Soyo) 5 Uo di x bc os we Bae ass OC end ee Lo Va eke oot 4 
influence of concentration UPON... . 2... beac eee aise bows 18 

BUR NCWANTMCIINCS IUCN 597 os tivo tk S'Atl ap Sins eco Bo wk eid ningeoe’ « Mak edee bale La 8, 10 

PRMEMIMETER Ar Sees ee en Neh eRe ee Wake ea oa wae end Cras Min ee 16 

EOS ce eck cis Py ee ke eet s SP VE Soe Cpu bruno 6 

PRCISNMDIY PRINELO A. 0. Se Wie lei goes Ces aie onda een eae 10, 145, 334 

SPIOEDOLUE DONTAODIONIAG. 6 Sa ws vhs 5 0s CRN okie lees 16 

MAANEE, CHRPE OAL O55 5653 9 0 Gk SiON woe Mala o kic ake & eee ra een 9 

MES 6 oO cceini c's Sateen d dee Conk Poarhww sg Me Vis Kine Oe TELA OE Seas 10 
TCS Si aie 05 cv tis Vale aU aves dbs od pea kedbaaean 17 
NM Shh Sai sg VAs oS eld Male Wei oe Pee ee 17 
gg Se ase sais) ie Alace aero 0a woe wd eg eae Coe Pea he See 171 
PiPe eee IS INTO PTOUPS 6 6 ons hoe ca ee gs we bee eve wh ales ews awous 284 
DN, IOC EPOITIN og CSc di ak pri at bo as eb @ Chal owe UN OD 70 

Re fain y csc Vales oe dees aches ee eclene plat nerd chs Bip 93, 102 
BEAM EE ASTRNCNIG BEL Pe ore one b edhe oc vba sabe h eoele d amwlemd net 90 
MUR his CHEE ME gS SSK Or5. c Vg wes oss ws a lnahwe Src obs lhe wate te 332 


HOIOOUTIO CHATS ON IONS, Fs dics ss 0's v.o'0.8 acactse Minions plasain-e a ad boy sials einen ahaa 7 
Electricity, transport Of......5..ccceccccescees evoke Gaecb Glaus, Ws aoe eels a ae 5 
BRCOP OG sina nies cog A owen Soochow et eh cm One Rea eee IR RE CO ea te 
Electrolysis, explanation of. 2.056... oss ee ccs 0 eke 6 0 a0't 5 ce Pelee Sian 
ADO UERS, 60.8. ccascla asked Wee eh hoy 6c hae Lice hee ae ee ee se 


Hlectrolytic dissociation... ..00..3050. Dc. Shae e does sae eek ane ts enn 

“Gnd ByGrolyels..: oss bc cceoes oe PP ‘ceed 
Hlectrolytic olution pressure. .o5.5555 5 jose caie ace pacaccceb cap oceee ‘othe oa i 
POCLEOINOIEOS BOTIOK. 352s Sebo bred ore cepa Sculls oon he ce ls ed 37, 41, 43 


MENS OOUEE CS STV aod wey dk we kk Eee REEL E Se CNT 8 pag 
Energy, free or available. {05.05 55 Say ds ase a eels oe os Pen 
SE Ce PRN int Dab ra tin ake erry ere ate oe Be Pa, Fh ‘6203 
BOREUITOD 50-5 Sic 23 Scr ag 8 40 e'eig & & ORK Se Cibie oie let ce Gn tuly yO escheat ann 
Equilibrium between a solid and a liquid.......6..... 0. ccc cece cent eee eee 

two liquids... 0.005 64.cs 5 os os eleee 


WOWIGPONCGOUS. <a. Ss wile bw Gk ote wea bn b.ccdve wy De eh 
EAPOTVRICTIOS 5 50 Sinn ohn a c's se Mek ae bas edd ebb. g des 5 vg Sah are Ol eee en 
PPOIGM FOACHONG. 6.425 PSS a aie bee ewe ee ke ee a eee aye 
CS 2 Fo Baers 0.4 4,6 aaa ORS ae We LM PRE E Tie ken evi deere aay 
ibabstrse Rea ooo as PG noes hob bdo bo bao Pee eb ee eH." 
MUGHYL BOCCRES, 6 o.oo Sin nce aie Bou woke Sa tea be baa eS hee pce ee es eee 

maarise Citalide.. 6 3. ise Gs vs ores bins wee oho che ss O0Oy kee 
Ethylene platinous chloride.................-. Nis» whe eS leleve, e hecy a 
OSIM Ts sak os wee oy) 5 5 ii¥s50 ood Rac are wa able gelato 9e eis theo ae 


TAVRDOTRUION OF ACIS 205.555 5 oon ale ped cei es cle» oe 6 0 'o'g mb sind 5 5 oil aie 
Pismenination, prelimiwmerys 62.66 esi vhs ols ded beck es ede ce ov cae ae 61, 423 


2 er ae Pee et eee ee eee PEPER EIN TROT ry 
Panne 6 BOMON ss 5.0 ipso 5 6 o5 65 ak do o's saws bc tw We ae a ee 
SERIE Ge ac ee Pee NERA Ree ee 78, 125 
SORTING Fo 55 Ve Reh aD iin vee ee Fs y BOOS 904 Wee we 


WOETICVANIC ACID. a; oa che Fake Ee bi inks eRe ie Ae ee eee 321, 327 
tt a oo a a Te ers Were eee ya a 
decomposition on ignition. ........ Ve ree 152; 322./ 3 
EMSC ITS CN 55-0 d 525 tiene fdaiade so! 05 Siw ee 0 Vita ha ine aN ot sc pe 318, 8327 
PRTORVODIAGS 5% 05.5844. ree GW gs 5908 5 oe bu Seok eae aed tea vende een Ss 
decomposition on ignition... ..........0ceceeccceeeeeeees 152, 320 

Perrone Compotnds igs .s oe ha cs cascade Meee rede Raed ices 6 bole 144 
OXIA DE s..ios sieve sataas base Rea , 29, 30, 31, 32, 148, 147. 





INDEX 525 
PAGE 
Ferrous oxide, detection in presence Of iron............. ccc eee e cece eee e eee 148 
sulfide, theory of dissolving.......... ee Raglan AUER pert Wistnl 145 
I ete Scr ele re no iss Wie-y ole.) Fe ee awe elk wee oe oa ES c Bee 58 
Nee eee car ee, Vale aie e'elay acy v's oPec ole whe wo Hale eg big Oala'ge welt 58 
Fischer, Emil, detection of hydrogen sulfide............ 0.0... cece cece ences 336 
I et Fo aaa gee plein gents vine Ovo eke vewee ges ormie 64, 427 
ETE S Sl cleft ear PaO TS aa crag pt ale oo tian ald o sree nis ea eae waa 62 
RG ee I See aide ee porate unre ne 6 61, 427 
NE ce eT haat, sdb er sas 0 Gor re RO ee hel dah 6 van Pein erga a aig alors 115 
NN Fo 22 0y giv lara rats bitchy e's pc @ acacels egies 8 WS + eset vee, 2 6le 410, 437 
ES oy SSE Omi eae Sire 8 ey ads veo date Weta ang om 407 
Se aNNINES ST AMINECNTELON BTL SUNIGALOS S = nL s oe ce renee ee eb bee nen eececenepaaeee 409 
PSS Grae ae ret nage 4' ate alo elated kak Be a a bate wie’ Kase 8 des 102, 406 
WOTIMG MOM. verges ccs ee PR aeare ae idee ard etal Me a NIWAS glee so hess tata a og aOR 311 
Fosterite...... Re eo ee Cog Mie as ih ne ee ober oes cxetnat saves eae 93 
IID or PCMAG Sa a pial aon eS eee ba oly py aetec a eb wwe 184 
Free alkali in presence of carbonates............... 0c eee cece ee eee ence ees 455 
RR re Sans We kee fae Vis ORTON REM oe Dae wikis ws be vee 2OL 
energy......... ES PRE RE RROD yee ie Sef ORO oS RC OL PE ERM ae IPE eve 140 
NE Ee arte eh cet cig ig glaigtr cate nV Ok wiaW staimr ele aietsce cs wn gee sae 121, 122 
Pe re Oke Cea Vico P RG hes) ebay OC Wedonk ee ee. 263 
Re) Re ener gi Te gS clea Eg Gi oe dio SMe one We opeheis 4 Geom . 279 
Rie ee ecm asie Cate Sirk sigh alsisle Sipe hae oon pb wie Sb oH 0 wwe 61, 63 
G 
I ate. eine Oy Net ate Gia d dese bowie y cutende wk Gea ws 465, 467, 478 
ET Be ere tell peg als sy Mey Sarg Verwle. dre k hid id ee Tag emia 0 ee NY 205 
eee ig cig Luo e pA WEN elder Aue wocig eS WU Ne easy h 8 ea UE OF Ute St 125 
REA Sere ee eS ONS SLES eR ues Dene v4 Law RE ee wae ee 172 
PeeAtION Ol WIUMINAGING. 6.6 Ss. a ees we oe views se ook ed oe ples 61 
Re ins Sn ie F< SG eas wee Woe 'n ¥ eme woke da end em alet sales 122 - 
EN ET tne Lao Seis aye ee a's aid ee EW A(e 0% 9 ROMER E OEE NS 310 
NED 5. S007 0 oct C77 5 sas be a oie hee see yen ss os eH ang Sate ees 1 
RED RTI ogi ton Saar gh wre ola eer ae we aise aoe WER ae teams 440 
See oe SS Ae e's Sac Le hw Uaard Ae cw tale wi areed. im Cote eee ee 172 
Ee CRs hg ond Sc Path inns VR aS ree en po we Re ese ol 461 
Gold, detection of small amounts in alloys and ores.............00.020e00e- 265 
EN es G0 Gh cS Sole 0 vie hast he Co CA Bogus a senasipe URS Ee ee 263 
yt ca foie Liga 6 SHA aE A Seale eee ES OOD. ros wooed 261 
ENC SHOTIMNIN 6 it Vie ook So eo ys els ie ca Cokie e he's 270 
ahr 64 oy ae, alain OD oo nea Fw ba OETKER Hae SOE SN abel 142 
TS Sonics Ee Le les e arssk'd oS ie Sa aks oe sag aa dR Meee 9 
aR a BS aT. on Sd Ski's: 3's a np loug os Sv CM A Vee mee ae bie ee ee 221 
Greiss, Peter, detection of nitrous acid... ... 2.2.6... 6. ccc cee eee eee ees 332 
Rees OP BOHN TGO Co iice. 6a Seale ws a oe a SG bec wiele he a cites eb ae os 284 
eae aE SRRMOE INES TINY G09 oy gk So ora 4 Gane. ose vie, co wUb Ie s Sie oY a Oe ee oes 70 
Guldberg and Waage, ‘‘Law of Chemical Mass-action”...................5. 15 
MRANC LN Mas TEIN noo cli eA e WN C0 wh. wb W's Aw 0'elb 00 Sep dw prostean Sees 238 


VDOT bai Oe Ry FORD 000.09, He ees bike Viele aes 2 ee F.9. OO "Flee G) OF8,.070, OS 102 






526 INDEX 
H 
PEE OK a ats ale g oe eK Basie aioamerg AS yee eee IN Are | 
Halogen compounds, insoluble... ............0.ceceeeeeeeeeeeeees iakeas er 
Halogens, detection of HCl, HBr, and HI in the presence of one another..... 306 
in presenoe of eyanide::.. i icdalieniiiea kak dan aemcelecbene ene 
ThiIGS RD ALS 5. hh ied eh oe Feet a nip lee eam . 
oxidizing action Of. 9 6. 6bs bee sue ae wles «8h eck oe 
TERUGUIORMUG. Oo Bee nnd Shes Nee ae, Hay eeee os OER aa Webi oe erie 
UDA, Ee Sain Eee SN MSE CES rep 
of formation of ‘water... 4.6 60 os 0 Sw Eee lp oo 
ES NG So Sega See Renee Date SE Ly ae Rae Aces gists 6 souace sa 'e-5/ aot een Le ; a 
EE ESR ie Ve RE SRR Se ees Rak RR aN te 351, 405, 427 
Homogeneous equilibrium. . 0... 00. essence kc eee os eb nos bie own wie Seen 
MORRIE WF ton aie da Riek Gey Liew Ho eae aisles Wa sack oo her 
RR CMPNE RR ease kin al carn ca eR a tbw vole ce CNG ah oe Ute eres Ee pcs many Sos ‘oy 
BRUM NICS a yet 5h so 50650 pu fw boy VERO RAE Tee ote oh Seng ew Gv 
DAMMSIES aos eS vib ne iin jared Labo Btatioiee ous ably peel «aon ys coer SON Tealh WA S c oP 
TIPU IND 5 6 oid 5 2.c ica we ssbe ase. chlale nex eine PS shh he at a I as 
MARINI FUCA 55 inl sip ev has eas rs RRA Sve ov eto hv eae eae Che bee ve wines 
detection in the presence of HCl and HBr.............. eens 
POGUCTHON DY. a. 6a) Yessy wise ioe, 0d ek aca pak ad ale Rene an Oe 
MESASPOUUORIING ACEC aos f'n eh a'y ond pare glee 00) Reader lott 6G a kh socal toe rede  ! ‘s 
detection in the presence of HCl and HI................. 306 
Piydrochlorie acid. .~. ii Seis ob cn a be Ss ped Op ae oe 285 
detection in the presence of HBr and HI................. ‘ 
HNO; and HClOs............. 3 
oxidation. by nitric acid: #0030. scs's see dels 2 3p ne 
POTOXICER: » 4)! 5.<va ey ks de sie ORE oa é 
Hydrochlorplatinic acid. See Chlorplatinic acid. | 
po iter ia fe nn: cei, aa RPE em mre eh : 
In presence of halogen and complex cyanide.............-- 
’ Hydroferricyanic acid. See Ferricyanic acid. ! 
Hydroferrocyanic acid. See Ferrocyanic acid. 
Ie SSM SEEALON ID MUBMAT 555 20. oo ch cp! 3-64.» 0: cg Soaiutin’s ehar'o apo. 9 Mee See eR bs ea eee 
Hydrofluosilicie acid. See Fluosilicic acid. a 
WIRE 2 oss ay 555 hsozs octes ov. a nk otc 0°) ae calls kb eee an 127 o@ 
Hydrogen ion, effects on hydrogen sulfide........ 0 ...... eee eeeeee 47, 145, 187 
Oe ROU her ee Siok, ue PAIL ue eee baes + 6 pls eae 28 
METOKIGE ee AS i sige Rais ache Bis nb ene ean ee ee oe ig 
ORIDRTIOM DY 565k waic ORD Ea hails eos 3 ee ‘ae >. 
reduction of permanganate... ...........0.cceeeeeees 85, 1714 
reduction by. .5 5.4.0. cs0s caved sles tals om seein 09 el nn 
sulfide........ i avapal ade ale ajak 4:0 ¥, oC hone aiee Se deeb bie "htace, een aaa 334 
as reducing agent... .. i. kane sins bales 5 oe ny ee 35 
detection in presence of H2SO; and H28,03..............-. 390 “oe 
effect of hydrogen ions on... ..... 2.0.25 01esdeoeleee eee Fae {A 
AOMIMALION OF Fis.) 5S os". Fa aes steele es a 10, 334 
theory of precipitation by 0.0... )5 50 o.e feds eee 47, 186 a 
BUR OIES oc ci05 5 ow SEK hig Oa ts, «oes nce aS HOt Mar HH nig a ee a, Se 
and electrolytic dissociation.............:0eseeeeeee ve +0 60d 4.0 en 


) 
~ Se 
& a 


PAGE 
Hydrolysis of ferric salts... 1.66... 0 sec eee cen ececeees Said ata lle dead ate ere os 148 
EE RAINURIS RUM ob ois 20 a alld 5.2 i ldwhe cae ide! avid. w peerace leo p.dimdewverd «Wee o 158 
PPPVENUON OF. sia se lek Alea beibie hia 6 Siake; ee RAN S o SLa'b9t 51 
MTN Mm FE Min eet Rec 9K a)! gan ein °b 20,4 dala pga a ha abe a Baers wis esviehage x A 127 
TNO Sc G20 AY foe ety cde Ua Wisin bac BU doa V obs en oe ov PHIL D wD ee 334 
Hydroxyl ions, detection in presence of carbonates. ..............0ce cess eee 455 
EE as We P00 wits ci 5-4 ig Aro Gig sluce hw bdo Sel OOS Ts ded « PO dwn c dhe 2 
Hypochlorite im presence of chlorate. 0:0. 6... cece eee eee eeececccces 397 
si aiy Linx ¢/os, esos dibs: (idhe bein Vo ha Side cares View denen ae. 294 
PPV DOCIOLOUS BIG 66. oe ce ce eee Gamera g sas ¢ aha, stele heap AY tae 294 
ENE eR IIIGY Of. Oo os Faas eens bccn cceedes cesuceceeteneas 345 
NE TLENGD ol og ao yok ok" oi ale pcs sv GAleie 6 slaves Fe wie ae bee eae us 345 
M86 WG va eiy sack wah ok kash cvdews sued cewen ep hink 487 
I 
a a ato 5s. 82 0 9 ela a Ale slo vibpepele, biace/Sa Wane Rat J wage’ 61 
I ets rcs nicer ssc, p-crake esac ee eae ese aw oie alec sass le Myhan we'd cod 157 
MeRRtRTMINS, TWHELOING ANG OXIGIC .-.). 0... iow co ee te hie cee elec encase cee 68, 427 
ELEM OR ee Lo ony og Rectal dey Merle a wiselp Bia cme» weiele ages Oe 56 
IRs re Sather s, Wiz ig Sot Ag eM ho Eve ee Be Oe oe als ok eeelek 54 
NEI en Oe oe A GC elsibtege Mice dolnel Smee 54.5 4g avn dle ain 331, 394, 510 
Te Og eo PI dats nsec ace cul eiarae Sm mie wince ea 0 deat 198 
SESE CORPO ON) C0 ET (0 |: a na ao 436 
ERE PAS FESS (eal AA Te MOA Ws AN | EE RE Se 434 
SRRMMUNERLGRE Sire csc cAI a Sabina a, oie ahaa, maemo plerete ig at 436 
euiraves, method Of attacking. 3.0). 2) fees cn ot bis on Wewleie oe lees 437 
EE at Ear ae ar Pa go 375 
TRIN aos oo.4-< = Glass. tc oss Por ed x WN rae ahi sour ey re kbar: {mal Stk Oo 8 375 
RIEL os Sag ee te ET SIAL cad wi dhk aia Vie wha ath ayes oR Oe ONS « 8 hevara piel « ath 306 
EE SCRIE seed Sa hie ch Ib edeie cau. k Fyoa he Ue REE Os ase HE iohie Alea ties 301 
Todime; detection in non-electrolytes. . 2... ek ee cet en ene’ 304 
FRCRCNOA OE OUNING Coa) oka Us. wks oo Meda avon eha en 306 
equilibrium between two solvents. ............. cece eee cece e neces 17 
I eg De are en arabe dg ca 'g & oh'vi ge amie Mie wise a, 256. 'g) blaine ow RW RT EER 304 
ME OE Ge gre an ys Ge oko 8S ao meee Hees, Lak o> been bight aA AS 305 
I crak tL Ge Ie rg tice a Vig bales, oo Rie stile a'n « pbaige a SRI w Wages 4, 10 
EM Re Sad Ce ahe oak wa SAG oka Saleh ee oe chKe RM OREO Oe ey a 
Nh OOo ida hak ems DAT Ga ee eo ee sll oeapret es eaab 24 
concentration proportional to conductivity................. 00. e eee eee 5 
TINY So ane oan aod al cuit cranes sea Rie sein of RE poy Aree Ses 5 
a ea eco d- ty SSS 5k ald ye ne Yee a 8c oid oa febnc Gua Rialems 10 
ROO io or A 27's oc ce CTE HEME Rais Coo ON os ote wewiew oN 4 
ee Ua ik yo ASG Tacs 2 vate ork kia we S Sean Ole tMtenen 26 
Ionization. See Electrolytic dissociation. 
a EIS SS eS eo Ni Nt ara Ie Fe ete we 53 
A OURNGOS Ii CONCENLTALION. 6.5 bc... cee cu cee nw sevepese ten 18 
SRMMMNE Sg AN Loe Crag MN gayi di ud Bete © ¢ o.c awd Ch SAS Wem eo we 19 
ITER OUENN eo gunners LIC Riad Spe hal sv ou di doa vole pape einceda ia aiaolehh aes 8 





a 
528 INDEX a 
‘ PAGE F 
Ionization of salts of weak acids or bases. ............ cesccccccccacccceses oa 
ei ee Oy Th mee APR MEN RMN OME EG 512 
EERO 85 Sins See's a a whe Wey 0k bbe lc GOW pe ke ates Te 504° <5 
pL OE OE Pe Pe ae eI merit eae PES Paha ne ST ~142°-49 
detection in the presence of ferrous oxide... ...........00ceceeeeeceee , 1488 
FUSTI Oi oon Fa KE sie aco en ee pnw Ol piace np ate 143 > 
separation from other members of Group IIT.................. 189, 192, 443 
state of oxidation in original substance: .............c.ceceeccecececun 4438 
J “% 
SVE EL OR EP EL PR ETE EET ORT PE EET ELE EERE es 413 a 
K 3 
Decca. s dal sos pao s 9 ow ca Me pwns vin aS Dye, ubcetdl asi ieee. Wig ee wits ona a en 125 
PGGATEDOENUG «065 62d as kvg 5 Soa-s bie oie SoRis,a Roe a, 8-Aie eptb ine Mpcdie win alesktes 5 lee tc 242, 
BOUOT TU 26 a5 aaah 0 0 2h-0 no Bis sk lv b's o4'0 9 lass 5 EIST carte ebb ca 93 
REMAADEOUS 5's 70s Ss eo Wise are ve Be wee wale vio Nie b Se Pak ae ee » Sheen or 
PERSMNIGE G2. hes ce ee S USED ROS WARK TO Seat Poles ¢ ob. eee en 499 
L : 
Laboratory reagents... os i: 6s s's is 5c ond oes 5% en cisely + fa ene ee ee ag 
CES ie 5 6s Scan G4, Vs bn eonTA p/Ble ete bila se, 8p mle thx 9 a altos os bane 55 
Pathan 50 ak Fe as oe 2 5 Sree eo ko hE ae pee ke ee 469, 472 
ARTA oo Ea keg soon 8 ie 6 le ceo are led nc ge “2 132 
Lead, acetate (sugar of lead)... ce ea ee We es 0 ee ee 341 
POBCUONG 53s boas vi ae Sk ok oe ae ae FG eR Pe wl ate 205 
separation from-Ag-and Hg"... ws. 2 .oicss vss » eoiee eee ee 282, 441 
ammonium sulfide group......2......5.0e2cceeeeee 4. 272 
: other metals of H2S group................ ee eS 272,440 
Buiate;, SOMMION Gio. a ios s a seeds bab nus.. uae dn yl eee ee 209, 437 4 
BS Oo GT GG Os MOREE ME OSE URE CAMRY 496 
DMG Fo oars OSes eka BRL ENGR O56 ee a oe ee 457, 458 
ME BANCOTROIILG 2 ists 5 6c 0 5 8% ‘oa ote eee cle 0's He lee ae wees | Sie ik ne 461 
RII 55 ane le wg ale oa tn hl G "oa ohm 6h aa lv'w bs nn. °0 aco se) 9 gk /bo ed ea ghee pe 154 
Oa a dR Oe LAAT 102 
TS TES Need oe A Pa SR eA a Ne ‘Sete 142 
MANS, ATG YES OE oii orace ss tals a tecorae lat shame h prs. soe NEE ale me) 452. 
PRRINMEMEG 525-5. ou Sneed Gos ce ve o-b btalarb ig 'asp/asueophasevaehtnvesesk a: Osea eleiw coleee eaten nana 206 
BAPTMED eee a os gbei arn wher a bletreal anda! wove SOUL Gao. aoe Sie eet ons lon fig a are ae 458 
separation from rubidium and ceesium.............0e cece eee tenes 459 
RUINS oan! 6: 6 Worn e so whee a: 0"0' wha wig" tw’ elsb'a Sovoret nse ls bia” 5 Tate es kal mi an ate eye 55 
RAMTE OIC aos \ 5's alba! no herarwly’ ole 4 bs ere i Sve-le' where ws aapily a pluie. Shee tsar aaie a np en 224 
M 
Bineriesin Comenit |. 5 i: o. ioc v.sv sic o's’ 0 vreln.e'e view bh wage. vei elalstele ee q:9seid 5 dle reen ann 94 
PUUEKUANTG o's fas 30160 10 4S +o bis aces a Noy Bat ad Bk oom eee 73, 378 
RPRMURGLS soos KS 4 harem Aowe Sis, nights a Sybele ee alae Se ae 93, 102 
Magnesium ammonium phosphate, solubility of ...........-.eceeeeeseees 75, 95 
hydroxide and ammonia... sccses cvs sueecoe€s seb oe seals ane 94 


PAGE 

TT RIAD Foe ioe adits Coe ess sides ise ve aeceassdwcarebevewe 93 

MRRINTENS WP UONE e712) cari ghes hated aos 0 'A/v a aaa ale a'a'e ove 0-0! 60h) ofp a niain oluhy 67 

MMPOLIONL IVORY GIRGIICR. 555. sa sac sce siclee 8 ad ewes eee asy > 96, 440 

Ns ye Pw ag calsh ¢ cd's ene COM NE alle t viud ou g bene Vers 93, 142 

ee Net ae Sek Sy. lic gla ale 6 al vate OU Ria Wek da Ae NO es ORS 214 

ee oda foc oe ON 6 dihc sb OSs avy so ee Sade vieu vCadvae 160 

separation from Ni, Co, and Zn..................20082 - 189, 192, 443 

SE aR oO a SP a ARG 169 

eee ON OS ae Nora age aC Ao. 0% 4 dolce bide ¢ oulaye s vie bois 160 

ED Foxe OEE Pagar Pat galwe Ee Oe eae eld wee demem giv anels 162 

Tg c Sra a 5 595) 4 eg 4 as ak k dog 42 eed eed hb Be 2 vole et 162, 163 

OEE EI rg Op nae ak aa ara e gare Bae ae a er 163 

tes Se eee U cig see he Gaede eke Uh as web weed Oey ees 102 

Re ee A ai! sw eal bea oN ep oe FA eh oe Meee Walbwe pap dee 142 

SET ELE LOL ite TER CY RET EEERCC Ee 232 

EN eis ine oe oi a. o's ON inay s)njalk o bik's aoe ee end ate es eee 13 

applied to hydrolysis of bismuth solutions.................. 211 

ORGAO ANG YOAUCHION. 70-5. 6 fs 6s Hs oe clos o eaate 44 

precipitation by ammonia.................+. 19, 94 

hydrogen sulfide.......... 47, 145, 187 

ee SSS Cs oa ATM Ce ae RE TOON 93 

EIRP TIE TSS ch ane RPO PSR len er 461 

Peeeereeann te OL COFLAIN MOEtAIS 6. je as bee eee cee renee ee ome 63 

OD LT 1. SE a ee GPa 157 

IIE rah, Sr ck Is Cad sa iw wid we ie Wwe Vie vee Oh 196 

RE EE SSS ne eg en eel ne nn a 314 

mercury, separation from other metals precipitated by H.S..... 272, 440 

ammonium sulfide group............... ¢. 212 

TIENT chy ke hah ny oats eT) As aia saab 201 

mercury, separation from Ag and Ph............6.. 2.005000 282, 441 

1 hy Ge a a ae el cook 5 w bane cw eb oe gate ape 195 

MIM AEB rey Ay cise hie eee RA RG Vek ck WRN ae 203 

PU UOTIAOMRISEANE EN 2 5 oar a> bce. ok: vilenaite’ ss hnale sesuce dw 09.0 ds Gan tein wae 204 

WIGGREMG ANG OXICIC INCTUSALIONS.. . ees ects vee ed cvevessceccseeece 68, 427 

Metalloids. See Anions. 

a EN EMIT INGO WTOUDE: >. 5 505 osc oe ble ee seb ucee ee ssnscdeameceveges 70 

MR re ie ER i Sy awh og a Oe wee walnaiaas 452 

general tables for the examination of.................cceceeveceees 440 

preparation of solution for analysis for. ...............-0+.000 430, 438 

ES ge ae Sieve te ne SEE AER TNE LR re tay: 77 

NA eC eee NA a es alesucea Walp ein es 33, 67 

Pen ein MMMM BOIUDINIGY (OE... 5 cack cs ce tues sence dee eee eee pescs ees 372 

a AEN, DP SLs case wi'sracg 4-0 Fed's Baw vA Re Oe 372 

na sarees pen weegidbieale 257, 435, 437 

ene Se isos ped ee bene eedoss oy 54, 284 

nist Ye foe ats ere Ny aM ces 2c oie 5s die 3.4 dae w a'bb tate wae oe 336 

RO ie EL rns vies bad «wate hanes Lane 78, 125 
Microcosmic salt. See Phosphorus, salt of. 

Mass, COUBCIM ME APOOIIG We ho ss a is hikaleie Gib hs chaboed wcbicedveeeaees 238 


BET ON eis Ty oe) ee RIND, sigue ES RLS, i a ee eae ae a 172 





530 INDEX 
BUNT GONE G0 Cavite gacdis Bs Heer Rore eit PEO vat TH rs va. 6S 
DR RGIGD, 555 oj So ac we a's 10 eee ERY Kd ane b aeons hae cele ene 102, 205, 
PORTING 0 aay, ib 40:0 jn wees in 6 edbellg clu lag baaoe an OI: ea ee te ei te a x 
Mitscherlich test for white phosphorus... ..............cc0cceccees wevevees O 
Mobiybdate'of ammonium. 212545505 (4. iis SS Wibieen tale a pale 72, 231, 379 Z 
Molybdenum, reactions.............. smi alls‘e/ ace one wo wae Ra oe ae . 491 
separation from As, Sb, and Sns. ss... a vo eww lee oe 4 493, 
Rintobdenite! 20.0 Me a ee a 491 - 
Mbly bic Reid seo ac ok eee Sek ne Whee calculi ae . 491 
PU DOGIOE x55 5 Rs .c  ajevegnssrn sin 0 -w wak Kia se le" Rel 6 aN ao aed a 465, 480 
ET 7 eS: Lk SRM MASE ae NERO RS <4 Oo 
TPAOUREBTRI TO ooo ood die saad 0 BWiniol gen dalbp a al, wos. bude. ote -atn ai'b, 9a us SSF a a 
DUN TEOO WIG se 550. G ce lo. Va ona bie So wane opus HTN oboe loses ly soe 28/8 Sate ok at 78, 125 
EGU Bacio <.creinin 0-5 0d B ass. eto sate th op cmd leg rue aaa nea bs Soc alts ey Pe 
Molal:toncentration. <2. 306.26. oak am wae > by Waleed a 4/6 Yi a ee hoe 
N ae 
eo a See Ea Pn POM ETON AORN 261, 499 © 
BYAGIIUG «onus sls 35 ops nice assess UEP DL Avy 0 bbe tyleteels Torok va ile ea 
PREIY Fr hs 5 ss < Sp Seth Sats Stl Le eco VER oe eae ee be, 
FOAUPNGHIIE, oo eal ca oe the Ti Meaty eta 4 toes» ach gh ane One aie ee 
Negative and positive elements. ............ 20.2 cece ee ene eeeees oe 
elements, examination for. . 6..0...03 0. 5 eee sa ve a 0) oh ! 
PRE STINET SF hd)p atk pion ere c SSe ee a a ede ap SOK iS oat ete a 469, 473 
Wornst Tommi s. 50. oP aig Pa ee sige v cgs ales Vo a pn ald eg |e 
Nessler’s reagent... ........00eceeees 5s er wslevanle aprplg pa 73 e Saibs tase en 73, 90, 198 _ 
TE aE ENE LO OTE RUT ye 90 
Neutralization of acids and bases... ...... 053 66 ccs) Stine be we ab else 6, 615 
eabiof. vc ca ah eo ew ebb ek mel aT ra eee eee eat : 
OWING Se cag old ¥ SG a wig a we eRe S ed eee pee . 604 | 
OE Le Se OPC Ree eC ba Mmr at Meme 172, 224 
Nickel, detection in cobalt salts... .........cccecececetpenswenceveces 176, 1774 
peelings. 6 Shs soa is 6a Sate k Bis hve 5.6 OR cs ieee eae 172 
separation from Mn, Co, and Zn... ...5 2.0... ese c esc e cece 189, 192, 443 
BS 5 22 FS RE GR wa Seste’ ow eipal a aNe}d a. tVagac ees apa ge eas. 174 
UOMO, 2b nk ais ole eb a aca ln wry 050 ew a jh a Telet con wha ypc tte ane aaa 480 © 
MOON 5 Vic 5.55 aba God coins hale vere aie lle der om Ses ee 480, 482 
separation from tantalum... 00.5 6....sseeeee see dete dee gegen 
Witrates, solubility Gf).6 6.60. le oes CaN TS Coe Cae Manis ele er 393 
Nitric acid, detection in the presence of nitrous acid... .....%'s35 +) See 395, 
detection in the presence of HCIO; and HCl................... . 398 4 
OXIMMHON Ua OW fs huis dae ov eS Rasy te hash eee ee ee ee 30 a 
he of hydrochloric acid... 6... 66444 ss ee see We He hes eR ' W874 
Poncbi) 2) 0 6 sid eas a:s00% ceels pedis var eee poke Han ee: Si) 
ER PA em i irae 7 eRe te oe), yes Pele alk» «ee .¢ one +00) 439 79 
Se Oe ere ea ae 330 a 
Nitrogen, detection in organic compounds, ..........0seecce cece eeeeeeneee 316 
WNitro-prisside test: oe kev oe oe ove eon 0:44 shinee ese sie me tie emma 60a 
MME, Cee eck ela eink cava dae gusd io accatenthien ae 330 
Nitrosoplatinic chloride: .)5 0325 4-.j.+ wis <0 sein se die'uin do biae to's os m8 a ing oe 269 


: INDEX 531 


PAGE 
aE SIMON TIRE MRI clots 54. eh aold caida SMSMAE a they amt oie easy ee oaks eae DP astna ack peas 287 
EIN «Sees tesa Pee st Oot cee cee éeed en pee eae eaisd a wakek 330 
Nitrous acid, detection in the presence of HNO3............. cece eee eee eees 395 
ANOS Sie a oe Paes F FTE Trio ORK ODE Ure Oe Reb Metal e Sine 329 
MIRROR Selo os oaks ad: rela ae Ce bial gia'e ov ed epeieia.e she unpdaaeees ered 10 
eg a tM get e Seale aig ey acy a’ ce als ¢) 40 vg ayiiene sca oka 3 
EN Se aI SB os dhe ghee ae a ya winly cob c'y ey ola wes TE eee nate 6 
Non-metallic substances, examination Of............6ccceecececceecceceees 423 
Normal solutions................ Le eeate alee Chae SERN od vals be lagi weet 9 
SS 5 iN ou ov Sw oa SRG y erelt Siu 00 ke bee e's 00-09 p wld Ko Woe 172 

O 

TAS 8s ee Coa sesc a 5 ed lo Wipe we oie Vpn Woah at ds Vow aa Vee 0b .b ete waite 1 

Es ae hd sss a bee bolas we abe oa ce bes Mh OS a Sig aia! k abot k 93 - 
I rere slat | NS Pe oe tn a-ulely © aye ocx) owls wh ouie, wn Soucek Raw BD 496 
‘Organic matter, removal of.:.................. 129, 152, 172; 236, 238, 240, 443 
“substances, tests for halogens, sulfur, etc.............. 291, 299, 304, 338 
IER oa gtd) ans sialyl vc aloe te: wr ploy has stg 's aha d's aie/s wy et re to ee 125 
BN a ie iso a's oan: sha’ be Peas acTen ai Rate tae Wigs Gees Nal ana weld al Bee aes Sar | 125 
RI eT Ns ear yo Vi k'v oh wig dln Bia) Wiper eb otwW. Ae Wy love’ oi chuya- wa Cade we 224 
i ee eel ge Bp Le arial hip re nsw BAe © aiden. wi hw blaisila Wa We #69 bideots 469 
eS oe Pr dS gue a yi plank ole Wik. eicle Cielvig as iefae Slubperes bawes 125 
IT ht Sas ee pn Soe lead Silay wR SS aged ok wpe ten s+ Maal 50S 
NL UF NR ara Vert sso Sab preci ai 3s iw Vee iowa gs Athy eae ae ead otis 504 
POOAMMAIIN, TEACTIONS. :. 2.56.8. eee ees Pa D a anwar Tat alcl Bee haw SES aOR Reb heels re 509 
EN Se Tene Mise GW Acd ily AiG: o Stald 0%, m:aShS: souces oo, BURGOS, a She eos ntoth 37 
EPIMMEEVION OFT MOTIBION «0.554 odo os clic ols wise vals oie a Uo ee alleles Da wa ely ge 361 
LEIA SRB Pa a pl Sea A Ieee in, a aneE tie, oh SMIeS ERE 8 360 
RIES Mee Sets eA Seer cic cae Aan oe as a aeie es Sk ioe =o oe ONS 360 
detection in Saelinbadiey Ee UNOTIMAL ROE Fs Suh aeradep s aighahss ois oe 424, 443 
SR MAAGIOT DY PHOETRATAN IE. 6S sass cv ols bcs alee we so cl biatecb ace 0 Grape Oe 171 
Oxidation, and reduction in vitreous fluxes....... 00.0... 00 cece eee ee cece 64, 425 
PERNT REDL a ime ga Lo tASS (ok alc ga’ aly pia eas deste pede cae 32, 138 
halogens........ Tee RD oh AIAN ct a ALP LEAT OBOE © Fr Be 30 
G0 iS kee eee 32 
NUN EMRE csc 9) Sata< ert iahe cna (able Vives kd widcbce Be cause Wea bata era rae 30 
MPUMNMIPRIATIS TROUT 22h 5705S. SIG oc DA Ras a Bee wTb ad 8.8 BR wie ee ale 31, 171 
OE AEE RS SLE NS Ss Sa AE aa A a a ORE CED ASE gy AL gt ae 27 
of one molecule at expense of another similar one................ 170 
EE DNCER EN aes Tee. 6.55 Sabon, Wie Sun Pacer ws om Riks ANA aba» Weetie Uie ONRTaa ON 287 
PIOMNMUB si lds vie ag sie kheee Yee oe ARE AER Tee 286 
SER hE res Aa Ske ats ay hla eo aan aceera Ties sa aah 37, 40, 48 
gn oy ie Ai chal tow cocoa Gao civ ¥ Bcwlamdre'n branes 68, 427 
I Cetra ed ee SOS NE wie g hikalalcis baph@awis PAICEMIC Es 86 

bi 

Me OSE ERS BO ai OR ny CA a a be WER ee ee 506 
Pea EI re ee Sr Re eee Ly vise «led gecslb ala aves dwibiwie gw ale abe 504 
Pm RUNNERS Fae 6 oc iy Soe oa ww bap awed (ae bake ewe 505 


532 INDEX 


Palladous ‘compounds sss... 65 6 fave wiess de deen eas NaW@e ba sagagie v0.9 ose >> ae 
PRCA Weis os oo bo wie cites fake owes be Sot Ca ea a Be 
PRPOATOORIO. BOI, © co sro: 6 ovo ooo Fe Nery widgets Wh? cote. a 3 eee . 

' Perchloric acid........... AN ee eee PP rae . 406 
Papen romnd aid so. sa5 0's os vo chidgs s Sok weeding ck a vee 132, 139 
Permanganates. See Permanganic acid. . . = 
Perea APmasine: MCMC Sos vam oie 4 core to ed cock tw kato a ae ee ee 
oxidation DY oy 6.25. 5.55 sa we elute Oe oleate eee 31, 171 


Pepe iRie BOI. Sie s 5c ok oa ne eee ee ess ip hekua soe a A 
PE MIONG ONE 6S Se hea etdig cab's Sidecar ges ie 488, 


a OSes a Rag kak oe Tia ees we Den ee 461 
Phewobpithedeha sla seals ae Secs wehbe eth tebe tuk) Laenel oe he 
Prorohates; Golubility:-Of. os a sy in Fh y seve sabe ook Pe slo hae ce : 
Pilospnides. aonition Of 4 o< i sissies Des ah rag Res woe ee wt 3s 42 
PUES fa nna eure hin uisien Sah Neo ac Ree 369 
Phosphoric acid, detection in preliminary examination..................... . 443 


RAPER. fe, c's. 3's 5.0 SAGA rhe pate e Dual os Bante er oe 382 
detection according to Blondlot-Dusart...................0000- 88. j x 
Mitacherliels)..c. crs oP can at gina ea 7 
in iron and steel. 6.0... 504 sats Le ea eee ee . 
pentachloride, dissociation of ........ Pic oar b g wakes See oh ane ari 
1 IE Ee Ne ARON Te iy TL ee 64, 380, 425 
Physical behavior of solutions.. ....05. 6.02000 4.200 iS ohess nd bee ‘a 
TENG Bee 58 hts ha ten ots BOP PG OY ed Se a ee pe 
Platinic chloride. See Chlorplatinic acid. 4 
ANNIE, TOOUONG. 6 ove Bock gn tow Cohn oid te emily Ninel alate ere RRR ee, 514, 516 7 
MONCUIONS 5 °5 35. 650 o's so eee paad wes wal Bas kcle ON ABs te 3 
sesidues, treatment Of) 006 5 50'S bv avs Hee 
goparation from: BOlGs 5. u65.4 5650 c esse ye bak aed Mos ON Ey 
APPS Ace ret i ore oar ee + Lececleds od Waa 6, catey te a 
IMMUNE AOA, oss. bas esc vie co Weeks Rew © Lalens © suk soos Geer een 


WE cs ko ose vc 0 ook bin, OWa bE wise Eee we ha ed ee Pe ee 156 
Positaye snd negative elements. ....5 <0. 6 ess sna be ees 0b nes awte ss 2 eee 
Potash PER 6 sho. 6 w'6 sign orb LU aA wre Wi.a 'o. « BRw Ano WSsgee Wekzt ol catwiie ta Ue 2 





INDEX 533 


PAGE 

POtabalM PersUliate. 2... 25. gcc ea ecesewcerecccccvas asdeéceVewsccesear GUE 
pyroantimonate............. Pee terior TOPE FP Pee 5 , ( 
SURMMTITRLON SUMOTL WILN es: bic seis eaics b.t.0.bsb cla aen'ae ve obs 130, 157, 438. 

IN eR OR ea abs CRE RE OO SOC NEOR a ehon a ee alee ee coer 78 

RCINMAES EEE OS oe, 5 ara’ Wal gad slateie vise id * caleeib hte ha tle eles 67 
RMN 8 seo a ig andl ol woth a.a°X «6 SNéip. 410 Gace be 01 Depa 37, 40, 43 
Es eet g Nae acu a(n ais e.ar'e sseie ast allele Sinca asedie «4 Mea cie Bw e 491, 494 
oe 0 a ea i See re Ae marae A Pa 469, 473 
NINN OE G92 eg od ahaa vs gk a's, W ofa ls aS ood Gedo a Mid Na din’s, Mace’ epee 58 
TT ORDLIMTALION «, v5). 0). 6, wv cpdsgied h odiels ae sa des co cieae Sddee a 423, 428 
ER nd ata sale GG cha een le Midis! eieyein Vosss-n) eA qs ste. a's oud ais bmw Rad: 0e 1 
ONT S 6 oo 5. inca 9) dia nik. hic.ere, sep nie, ec gersie'e eas Pde hate eo ae ty (eee 60 
Es Seas asco pe en Fas Wie daa FS ae Nl 8 pe IT whee hike ni od ele aie Qinvee's 224, 278 
ee orn eS ay ao g bk wae ales ele ose so tan, Led ee ee tate 150 
I SRT WOMOIW 6 oS So ohn a 04 210s Geo es Oe Wee islae ole naw walbeete 318 
0 SRE ges ie CADE UR RPO Be Meer coh eae Dotan aa PE Pt 321 

TIED er aed opt S als Ou. b eo oS aleve edearhi as advan Ne cekiods'e walk 310 
RRS ery See EF sca City enc Fda sate wie Ok ovlele nee veel be 263, 265 
aa ite ey ayn aisle suai ed ww aN oe wus ck ceeiee Rs pee eS 242, 278 
I ee eS wigs hii s aio os sine Ge 'els aceed dyes Oe MMs pia eca ar 142, 336 
LS ol sit ig ch cre ah as ee slSycleis Wael ea cate es ee Kwa eee omneh oes 465 
EE edo co 5 Gi > fae aieie, cree ae, ¢ WLS owe sin Shea. ace Pele Riadeeew ao ogc 160 
RESO a (OR ea ik on Reg ae veal) Lhe ped Rome ‘.. 102, 205 
RI SIPIBILUY OD 5c oo. aaa ok a0 she os, Fan aa 9.0 nye" PM Be Rie -eieigt oss De 374 
SC MNCINIO’QCIN, TORCUONS . ... 34.0 cies cs Sc bce ecru Sn ee oa Selec oiee ovis ye bie 374 
Pyrosulfate of potassium, fusion with. ................ cece ewww eee 130, 157, 4388 
Pyroxene-amphibole group of minerals. ........... 2. cece cece e eee e ee eeees 93 

Q 
TIMED. oa a) ong. 5s 55 cay, so ooaes & vehi os Mia news BA Rho DE we aiale eg ibe eats 1 
ENS MATRUUNIS Seo aks os alg lane sie aid a Beard ee dra bo eee Wee: Wo ws wah a ssbrald tree’ é 1 
Ie Sete alot 5a F Siatele gin cca ih ¥ ied amie w'o:PeTatg BIR a tue Rien « Hivigrace eels 413 
| R | 

LOT NE ook. soc vee Put Glen, bw oes leo dea w ON 461, 472, 474 
a ARE el SE oea a is or a Sei) | ag cae nee wy Seam ANSE AN eel Ra 456 
MEMS ICONS tik et POR Seo ores ba ose Bead ewes ob es 490 
ERR ORES OR See MCSE AR TA i alae SR Re 1 
EE rte as 5 Poe SA o ey dS. ¢ dew EUs ww ee Pele 26, 51 

Rey IEE Se Ae ee a WE Sr eae eat oon SO Re ees OP nae oe Se 26 
IE IC SCO 8g fice Fc lg ain Bs bie kus ie a vieie Rely Coie athe nies 1, 61 
NNN ciclo wtacalas eke FINES CNS 6 Vio F Rip OTN. Whew Oak ple Ghd Sales 1 

MEUM CD eS SL pt 5 Geek Pole eal hv pod Gos © ohh dle Bx aiarwhiss 14 

Se NERMNC ANG a og le es aer te taty np WEG a8 PLGRS a PEE NIEE OR RS Laie alee 75 

Mee TOF CUAILALIVO GNALYBIG: 33.5 055 0 rs W esi cele ca bee ha ane de 1 

RIN ots). yeh cr cca g Rt pla Ware yl agteievale<g « vas Ue ee es bile aps Selene aa 1 
er  IMNTICEACIOU OE 5.14. 255 Vc Cals PAs pose eb ins FON MeE Saat w ews whieh ues 71 
ICES 2g a8 s GNOME, Ging SA BR Rte mah db Rat ce ores Whee eile Nie wie es 224 
NRE kr LS Ue cats vow hes > VET ss oN Ree gO SINER CERO SS ae sc 207 


534 INDEX 








Redueton by Devan alloy... FOE: Ta albe La gnd aoe eke vent 
hydriodie acid.. sev bee cubes be wh gleae po leery ie se 
Rydtegen*paroxiie: ets Sich pred SME Ty PRR ete ble I eee ody an 

sulfide....... <i pclae pints Phase Ola! sectal'e HRS LW URS 4k Reon 
metala syria: S250} kb en MIA A WERE Pee yr TIS er 
nascent hydrogens +5) estes s hc ces vey ele ee eee Beg 
stannous chloride..... Se TRU Te eat bas OAS Gal oo vs eh ae 
SUN UGG BOG LSS sass vs Seclew Seen ae oe ahs vevsteebl ss) OOpm——— 

in a glass tube....... wh ATs edt yo ae eon a f 
upper reducing flame...................0.68. ae e's pc ca ema a 

om chareonla gk. voces Nak es can int yews eg ee eee oe Oe 
charcoal stick....... Dingle ath heels tote OR ces Wee . é ain 0 

with sodium, potassium, and magnesium.................. ‘au 


Reinsch arsenic test...... Cah hechlasy Naa. ey pare ove ee 
PRS VOESHMG FOACTIONN, 6.5.55) hie oss = 510 tno ole ig a in 60's wail One +966 en 
Rhodium. .... bP S SURAT SRG Suhel kean ee at See ee ame +i, oo 


Rinmann’s green. . ee eet y Te eet as Us ass he tee ree ee oa ot 
pietie wats 0265.0 <3 c3200: eis eet MPN eat ea iS oa 
PKI ACS. oo, PLRe ei. Ghote aL E Uk ee aen easy, heen . eee 
sah VORUEONE: i205 sig 84 5, ald is heals wiehn e uloup the Peale sha.» Oe ve ee hee 
My & separation. from lithium: and: cesium: 03.5 5..0.5 va sapeie eee ae 
Rub Si Ta 5 ye Rie bales Bk o WMD CS We NEE De 2 1k 
Peete AGO 6 oa. ae Ra CRE A Ka) Rnd R eee eee ee ee 
MEARE GTLAPENTY THACEIONIB. os 56's vg ote so, ce piglets vim boo donk ls retails 2 iat eR wo. oe 
MA TAs kha. é ole sed sk aS teh caba ayes eee tas ae Dec Gere ..» 157, 1605 


. 
- ‘ 


PE CMMISORIOVING 5 Sia. Cue Os yeas dena ake SUNT E a $ ple einstellen 6 alana 
Salt of phosphorus beads............ As Ee We hanglegig an a eee ae ....- 64, 380, 425) 
solutions, concentrations of................. da ba RR oe WN 00 ob oot a 

PRM s 4s acs 0 pig8 OSs Oe ang a ice PAE ha eeld che ia ea es cca Ba 
A MMRAPONU EID 000 55 5-2!2 sio:a Sv ovo ves Kw ecote kt how oom 0.9.8 10Ny 8 = 0 ld ore ek 
RRA TAOU OL 5 i:k vb nd sss EC OR Oe SAW TS Ca ee oie glee ene 
solubility table of......... Pst Pl Shes VHeeee ee Ga dione sv SIE 5.9 cae 
RIE cho cc w deine sc ews wid ean Sd Fede dedee Lube ogee ee oe Sea ~ as LORS 
<a aa rape ire nner eNerrarmeny Pirate y reser eS) wih F i oey ene 
RMUTHEOD GOMITIONG. 6. bis 6 diel ide de dedave 46 daithcle FOBT) Ne EV Are J a 
RN hse) ano> soo Aa.0\l 2 dbut Ad mod. cnt eomn weet Bei ss.atoaktea (1. 
Selenic SD; TOR OGIODS oa ol.ol's Wasson ctw ey mp tiieivin dl clapath eae Orne a 
Selenious acid, reactions. Cee Wiad tei aie ed CALA eS YBa a” eer. 
MIL iris cicsa.+ « lee Seas ere Wrculct sep iee Chee Ladenen bags tn 
iid tAtbarinints $a ope eel ean sud ed-caelad vlc ater 1 DS OL 
test for in sulfuric acid........ Ld inte Meee ek EER 
EE POETS TLIO IS TF Ly PP hikideneniad ee 
Sensitiveness of reactions........... daw eatd lidvus.d ee Ree PS ee 
Separation of the acids . 6 oie. pais vee cating day ewes oe vee ee eit etal ales 446, 4 49 
COTitS MCCAINS) 630s avsivi nel slcw ee ales grrr te eeeeeerees & 
five metal groups from one another..............++00+000+ 


“ . PAGE 
Separation of gold from platinum... ........-.:e cece eee cece eee eee e nes 270 
MCR RATOUN Ee cua ak cs Chea eet Cra Aes ess eats he 282, 441 
GROUD LES AE oss Ne iia aise Se Pe ee: 272, 441 
eT LN Ae ORE AES ger Bai beret h aay Ba 189, 192, 443, 444 
RRIRTE IN yo 'a sian é cakes wie es eww descr h Se ake 111, 113, 440 
COI Ms feo re Sete tay tag hel Mane aia Sie Os 97, 99, 441 
MATT APIEIIAY THOUAIS 57.5 5 oe cele 4 i nnece Co dine vee Ou oe niee baa lo’ 514 
coy Phi oh 5 pled lore wae a'e aia e ye Mad dal tae and 2 y tege e 93 
TOO Cera el yon acd ig Oly ain alaluratain: eet bos ig We > bie Sch Bb. ib Os 102, 142 
EE 2 SRS SE ear eae evel stake glove late’e “aed alas aw eyed oe aie 415 
ERM EMLT ACID EIS LUROR AT. 6 Gis gicae kote clus 4s foea ceed sve e ep ba es 417 
WOFRIY ER es es oii eee Ca ta a eae a the Vado ale Bh 419 
I EEO Ste IY) Pee wd aetna diac Draka ald Belt ms. eS wd horn. a0 8 416 
MEOEEIC IIMA D IG IXY BOOS oso es o's ce Sials vid ad atele Dara gtece wee tu © brawls 417 
EP CIE 5), ahh hatec wits lel aa Bartle Cae aatalee al DEA Ged wee 415 
pUOTAGIC’. 6... Ss. 5 AES OPE Sh UNM as TBs PE Rese Si YR Eee ea a ae 413 
I a Sash cn Raced fi a lh conteio dle’? acetate; Abie = "9'e-9je a te 8 os 420 
emmmner ita 18) OR ONG GUCCI: Soc. cies Wiha ie 3s oa whe Seay bo a bas ghele's 420 
Silver chloride, solution of.............. Pee SEN aM er RATE CW -cke oly: el RD 20 
IMT Tht ek CL ise GMa su hes Tac he aie a tee bp MgiNeS dpe 028,850 e's 24 
hcg fering G:C pines Lee GBR ALOR sb slp CHER awe Oba pint 449 
re GUA ens torn bgt tN os ypu a eA oN goin ws 278 
separation from lead and mercury... ......... 6.00.0 cece eee eeees 282, 441 
MI STC ai gta Weng She Oren ees hua y GATE Sand <ihcd boas Re STR ee Bae <a 499 
I ke OT SRT ae 2 aad en ek anew bile divas 6 C4 419, 426 
I tects So a 8 pk gia inka 'g ralnrge hie tM acdc s prstew RE RE es ee eee 178 
RE es gt nce SK era acc Goh asa uei des bie Cece wad LE wwe bs 178, 224 
IIE es ie ce te eat od ile eg ea pe week 6 Wotan aie a ees <aieik y 102, 184 
Ne ar, fe ee ee See Oe lS gs he ae Re ete 82 
Sodium, carbonate, ionization of.................00.2 eee ae Dette onus 9 
cobaltic nitrite, reagent............. WOES N es % x ACL RA Ee 74, 81 
RUE OG NG NUE y cde tg hse nN Gea ASS Bim ako. CIEE EAS OR lone 336 
ig RO BEIM Sr GE PRS. EARP he Ins Paste Gd, 2 Os Pa eM ON 83 
EI SA Got a SPAN Lice iak Reh bp Meets pate: ARR SE 82 
I MCMENUME EE, OFA ve at))s SG Clalsgl ved Va cu alayar'gl Mis, wkd pha tage Wk ewe ohe bw Ate 67 
IO ce ne eG bie be PS ECR eye Sede Hens RH e weet dye we 8 220 
merapautics and. solubility products... 60.05. 06. cc ee wee ee sbincecens 19, 23 
SOHO VOR Ee icv oe Ora ba eo Ae Re We oe lacneere Cas 432 . 
PmErGe TOM -MNObALIC SUDStANICES: ss... ede.c cose s tenes ccvs sis caccevuccese 423 
PUMEMNIIMOE AMEX SUDSTBNCO.: 6s crn osc ccs ceeds dese sdiscchevtdeseects 430, 488, 447 
RETA MULTI UIUOEE oa 55. bovy ei ealciey were ee cain ene Hoon 435, 4386 
ET on oo 4 vv 55 8 Re Kos + 00 ie nsec iol Wubi 5 4 be Be oo keen ae ee 37. 
PERM a rhe Lice 8 s-voncinin's acy wb ip hast paren ach alee arae es 38 
NE Sig Sa ut hare sabres cot iia outed Flee EOS Rede Cb eevee ue 452 
chemical, electrical and physical behavior of...............00000 3 
EERIE A SS ite 9) cepa, Nina @ eRe Sean igre hee car ey ali 58 
saturated and super-saturated.... 0.2260. cece cece cece ee uceeene 12, 37 
re Re Pt 0.0, eke IL wos We gates ae OA wn’ ov vw pn bi iele Eo a Wa leak aioe 1 
I IE Nt rk hg Saree yd wf Xe A ERE 3 eR OA EMIT SL eam 120 


Special reagents......... Oe EMT NRSC | AR ANE Ah ECE tre ER a eT 73 





536 INDEX 
Spaptnogranh. Saas disiv eps th Chia Cease eaters ok a 
PHMMTROUIOUOR ge Sad wed vin whee 8 > dhs detent Ale a oR eats eat eee a 
Spectroscope.......... "Sos Relea oe a 5 ure Se eee a 
Spectrogeapio analysis... Soy. oe os a ee See ee ee a | 
Spectrum, measuring the lines of............... Aa 4-9 Gt eee os ia um ' 
SpMAlATiLO heck eA le cok 244d Ga LR eae ae ee ne ee ia 
Spinel... kee ee eee eta eee ova nb Sea a Sead. are teat iy ae ” 
Bpadwmens .. 25.5 spose Nadas ces eto os ecb es aekeaeaee ibaa ee . 4 
FSCRUETIAD OID oS Merde 6 vince os Whe cee ehow Se ktee ee 5 eee Y 
ORMPIOUIGR 25 ois = og oaks wee o's bd Kips See wane RIA ga ee 25 
b-Stannic compounds, . 5.62 (22)... 5 boas sce nce hee 0k ben bob glpse lene 
Stannous chloride, reduction with........... ida ne tae ee 2 
DORNDOWLEIE : 2iits5 soe Paes ts dae os Pea er 251 
SGGTINYL CHINTIGS 3.025 ik SSG eye bi ee be atk ee oka + cow eueel peg eee ea . 257 
Starch and potassium iodide reagent. .............c cnc cevscecpeveceushes ot 
Steatife............9 ELEM Pie ee ei SPER PEO TT Le ee 
CLT SS es Seer ee ee ere Rema Me PERN ECGS WS 
SME oa wick 6 Ba 0 kee RE iis wlan alps a abu agit way Se ee een 205, 494 
Sttength of reagents. ... 4. coc. oes cs gee ad cde dude Sued at bwolie Loa ee 7 
PRGNRAANUING. <> a8. v'Aas shes aes cw bes poy Nie oR 102, 106 
ieontine, POMOC he lad a o-ddiain set ene Se bs ADIN ie loon ae .. 
deparation from, Ba and Ca. i 6.30.6 cee ok oe a ee 110, 440 
Gulkates: solubility of....... sae ara.a eal hacbrn,a ws 4 LD Sasa ale cadre 404 
Sulfide, detection in presence of sulfite and thiosulfite....................00- 390 
Precimitation, thoOry Of. .s.4 9 cee ces ous coin ks caer meee ae 47, 186 
Suifides, behavior on ignition... 2... 00.5 Sede ss cess e ce kke ars ao diene 3 
mee by Of. 2. ek td So eatin pee 3 5, VP eR ecules in Ae el 335, 
Sulfites of alkaline earths, solubility of..................... eek 5 ee 348, 391 
Sulfocyanates. See Thiocyanates. a 
Sulfocyanic acid. See Thiocyanic acid. a 
PME tied on Ss 5s Sum aioe ch oo oh hae, 4a het ak fens a 339 
dioxide, preparation of......... GA's Cede Wace aad aap Ae le oe 347 
Sulfuretted hydrogen. See Hydrosulfuric acid. mo 
Sulfuric acid, ionization of... ........5....0.. 00 wi eed de bss «alone 8, 10 
TOAGIIONE oo le eS scarhe Sa ee PE Oa RR ata Nee ae om ae } 
tests for selenivm in. ..6 040.000. as Ges ss arkalcd x ieee ee 
Sulfurous acid, detection in the presente of HS.03.and HeS .. 2.24 se oss on nee 
TOACHONE <x. Ci anc « © Bao ale Balto gaa Bae Bika phn ae 
Teduction DY 64. 55 0 34.0 d4 oe be tin u wrolete oath slate pea iets tate 
Supermaairated solution... asses asle saws Anbae Oe eg ces bebe 
PN a ooo ha 03.0.0 0 0.0 bye Hb Ow 5 See es gE Fa a aE ee 261, 499 
RII ns vb.o son te ua se Mink wg oa Ue Cee oe narele a58 6am and amen tine aaa 
SSAA OS ey rere hs COTE aa 
cL 
Ae oad a) su on 'dvin ne bc bb 0h 5 we A eT ceR lye We oie Oactre abe Deen ey 
(ee Le Sees SU te es er 
MPM FORCLIONG, «6. . 5-5 0.5 EE) Foe € 99 80 bev Re ise kee a ae nr 
separation from nicbiim: ) . 3. 266.555 eeu va ose se cee hee ae 48 
NO oa oso 8 Tas 5 cha meee wk Re PL Gnane ieee 244, 362 


h 


: a 


PAGE 
NREL GEEE 2 ise ios e's s cee das oo wae Nise Ciclo EAL AT AIOE Tt EES 362 
artrates, solubility of..............s.0005- OR CAR ip contd SAT ea GARI eae pak 362 
Taste, sense of...... eee gra a Milas dee 416 GR TEAR Ts SWRA TES wars we eee where 2 
NOL MTU CoE, sag GN a pets ao heb b-Aie wes Go's! 0 vishe igiaw Fie Ge pb 520 
EE Se Pha as Cae Bac gag in git bie pam Bi XSiriaiphin A o-ciy 43 tgs wie GRIP Ro Ie we ® 499 
5 RE DRONITIVL SER MUNN 01 ile a0 cae dino hee Gale G lg sew Sie dog w eeine OL oe 502 
I SORCUIONIA Ss ganic dos wie lane a. © sted <sih'o wpa wien #0 6 4b os soi B hiece es 501 
EDS eat Ai Ne Gf eis Gls, 5's bls caw wk ce 2 np Oe 9, 8 hoe CGO HIN 0 ye BN 486 
Thallium, reactions...............-+- fl Cal Pe rae A Aaa BER OREN, Sle SR 485 
REIN, 00 ho Le args 2. ola Bb insm ee Rae $Y dpe + Lidia Npleibloe Cave wb alps 485 
NE SR ec he Ae Day be sae bhaey sel tieeeece 131 
Thermonatrite.......... Preece Kew TER PER aaa iets ie alamo ie to 82 
SER ste el gS Ste 8 a gs ig gin nt a acs oes & Klas WEED Nivree 0 tn 323 
in presence of ferrocyanide and ferricyanide................ 327 
SERIE ID LS CVE ose o olacaks bone $14 dbcte ora sions ue bits Gee bie elelp¥ ve alo as 387 
Thiosulfuric acid, in presence of sulfurous acid and hydrogen sulfide.......... 390 
RORUTIOU. Lies Cie hee os on cs SN Oe ee te eee 387 
Ree a ere ity on Sy aake pa uaa Veneer amass 36 45 465 
RES Re TEA Be ln ee a ice on OA Ca Oar 465 
Rl Tee tae aon ce Gry Wat eMidia Dx ale Pils gin! oe wets Ee Swchld w sia sis 465 
NT ria ko a ay fh secly phere GR Ae VIR UN'S a 'ee So) Soa es Lest WSl ole w ae 496 
ce Ry pPre Le re tig chy Wk apg ah WED. Jlabe' ad <i Zip diab @ wis 4g oe 6 249 
Separation from copper Zroup.. .... 5.6... s 1c cee cee eseee 272, 275, 441 
Ne og ae x Vis hg a 0 c.p ei ace atv ao ype ba daw Se 82, 357 
5 eS i bc ienajy, sc e's S's 7s.a.v on SS ay Shae a Riri ee Pe 249, 260, 437, 462 
Se Pott ca ig SP cs he oe STA yan, 4:8 ipl aagyid Lies be pie CORN ole «eb be 8 157 
Titanium dioxide, analysis of.............. ee avis cst Gish TOR WAP aie Bs ate 438 
MMMM EER Mire. AE Peet cse aie Sik se Le ENG «  cdipial wiece 157 
SS Page oe a Ae wh Ey gk Grae tale 2 aces Wand Sd ME Gh Bee aes des 0 8 125 
, ENN Was eee ks ald 4 ie oes Ww s¢ Ceo PIG ans Sass 125 
MIPECTISG OL... 55s ccc cee ecw anes LE RIE: cc Bi” PO ey ie EE 2 
Lo eh Ay oy eid y coy alice Bow oidclghe Mee e% ¥ 0 vein Saieih's 93 
IN eerste gies LS bre dio ie sk a eek s Wald tawe sy, xs MOIS Sarees CLK DOs eRe es 82 
CS 5 88 clara Ga tha acgp 90s Ps pe ROA bon bv ee Ud we PP oe alee e 494 
ST acre 2 ea iG geo SI wer SONG, Ge in I ha a Me yb Loa lates 494 
NI 3565-2 snes en lacgod Sree p'a bls EM ae sad Me OU ah 55, 74, 358, 464 
Or ay 2, ab are Bk SIaw Nie kd eee ovine s Boat bNG hae helee eee eee s+ | 147 
Re sc oot Sig peleodin wb eee ON es C3 pm ONE ce weld eee Ohne Corea’ 125 
U 
a oS ios oes 6s. ood Pe da wee sic ces nee abe ew enweh eens 172 
oir inc ee ae atc ks Cen asda oe soled Veweehdeuunebpecse 118 
Ee Pr kan eens kG Gg indy s ek An ole va py oard@ a reno dueone se 118 
a SR URE Ag Rr RT PORE Fw,” eRe et Ret et Oona 154 
ihe be eet esata ae aka eresi'o. ew 2 Who WE AGE GOL Bea we 154 
Senurevion from Fe, Al, and Cr. . ici. dee ch ve en 189, 443, 445 
I ra oe ate nats Ae rg teas «Tete «aR k oe Bein clan oem 154 
I IIMRORE 00 Toh ciy ook ec ve eR eae ORR 8 OS 87, 333, 396 
NCL WECM PSOTIG IN 2 i S6k ace Fics bev av es ve Ske ss ua e's woe ene cus ees 238 


RTE: TCTOURY Bi Oo cd's Fetes Rae's 6 wate a ae e's ve ee adepees 203 


5o2 







: - { , 
538 - INDEX 
oP: os 
WMMEIDME IS soy cca eae ee eee ee Myce uape pad sea be Geeeahy 5 a 
Vanadie acid..... Lae wire e's uhh Robo aed hicks ides dhe panna des vahe eA apes «el 
Vanadinite....... A POE ee are pre A TMi tren AL Say go 102, 205, 
OMARION =, dean Wn Viera sip Rebar aie a eictents igs 48s ASHEN MeES EE ET ait mA 
. detection ‘Mm: focks 5 56s See os sie ee oe ws ed elec nea 
VARA PBIB eo cee ikea Sp Oe aad ace ee ete 
WEMERIGO oo c aioio. oie ora pip. sib a sales lovee le Muse ecg aus ole bia Rite ak ha See 
Volatele substanoess os. 2s sisi. csyepecaicle: 6b \a!p a 3ladueipte'a aie eo 0a mR a eee end 
Volatility testi. esc. e e cere int 80h an emlece dap hieke Tan > Piece Jean 
W 
Waane. law of minas-aetion. ...)) 0 ico chic e'sa ss,s eae 3 9 1a eee aie 
Weall-pavers, ATSONIC AN. 1.6. bo vies pas ep ebeie se 5 05,6 Wale ema eee 
Walker’s nomenclature of ions............... Se situ « + scp} 0.0 etn RO vee 
Washing precipitates 24.5542 shee eats rae ete * stiecs hw each ave a . eee 
Winter “smmonia in: os d,s CoS as s eee evn co eos okey Cee eee 43 
free from ammonia... 2300655055 cds cnntensryy atone: See ae 
GEIS Sis oiece pH A ale laieg: Oe wha Re o Cs dich baba oe ae 4d 
heat of formation from ions.............0....00e00: (ius uae on 
MRRPARIOR OE sass et a sleeve eos Waa Ue ey eh wae lle Suge pace 
MIN Se hs ae ao SETS SOON ee ae Me att ae ae 
SW ereaenvartdae: 2 ooo Se ee os LAS EEK Lh aA cee ode AOMORI Oe Pe ee ee 
Wave-length spectroscope. 2... 2.0605... ene bee ee pees one a ens cle eee nn 
Weak acids and weak bases, salts of................. eee 
Wet way, reactions in the..... gift sigarete csc RG ca 4 b,8 7 Rm aoe ata a - 
TL oO ieee rR SPORE A TIMES ARPA CM Nae IN ee RE 2, 10 
WOhienthes ss Sa a eee ices» ive Cl Fa ieee ee 
WN RING cs sce ocak oe eu bo biel wine web elie ccd aie con ape ip oe le a 
rem cc Sy he Ste Ro a ge er 205, 491 
PVR MARLAIRIEL Gig ox. 6c 555.6 an gle elo 5 5S BU oh cotw dec sient’ ete ne ee Its ale ae 
x 2 
Doe ES ED A SORA ee ea OP i ge Re Mae OEP MET oe Pree . + 00kee 5 Senn 
\'g 
RENE 54.605 vib tb aalaud HOPS Ae Ayah ean eb a vee ko iene ope 
eameintAMRth:. 35 6 oooh ca doen, tw ce oa es ee 467, 
TRUS DOUAAOTIB, 2. 6.5 so 60S ive sa nile ave stale Ghee a 0a wivtelle Site ol ae eae 
separation from Mn, Ni, Co, and Zn..................5 189, 192, 443, 
PRIMA eas. 5 os, bs iged HGRA bib coe M Ee Ate RIR LL EE ODER Oe eT Re 
Mite made, precipitation Of; oo u6 os d es wd se de bel a AF ene Bip el sae’ 
DRI y ciusncs 0.2 0 d's.0 0c erwla’s o)m Weintelu yt ab pie Wel aie wre ey Nie ase me and a a , 











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